Skip to main content

Full text of "Inorganic Chemistry"

See other formats

CD = 

Sj<OU 160459 >m 

> C "" xD ^ 

^ DO -< < 

z ; > 

OUP 408 16-644 5,000. 


CsS^No. ^CJ (, Accession 

Author H15"J- 


This bo6k shouljBbe returned on or before theflite last marked b 

U vy i. 






E. J. HOLMYARD, M.A.; M.Sc. } D.Litt., F.I.C 

Late Scholar and Research Student of Sidney Sussex College, Cam- 
bridge; formerly Sixth Form Master (Science) at Marlborough 
College ; Head of the Science Department, Clifton College ; Member 
of the Royal Asiatic Society and the Mediaeval Academy of America ; 
Membre Correspondant du Comit6 International d'Histoire des 
Sciences; Examiner in Chemistry (Higher Certificate) to the 
Universities of Bristol, Manchester, Leeds, Sheffield, Liverpool 
and Birmingham 


First Published 1922 

Reprinted 1923 (tune], 1924. 19-5, 

1926, 1927, 192$, 1929, 1930 
Second Edition 1931 
Reprinted 19^, igtf, 1937, 1940, 1942, 

1943> 1945. 1947, 

Printed in Great Britain by 
Butler & Tanner Ltd. Frome and London 



Science has now been taught for many years in the bulk 
of the schools of this country, and for nearly a generation 
it has been an integral part of the education which the Board 
of Education approves. The somewhat paradoxical result 
has been that there is perhaps less respect for the scientist 
than in the days of Huxley and Tyndall, when science had 
no footing in the schools at all. Few would be bold enough 
to hazard the assertion that there are signs in any class of 
British society of the scientific habit of mind. 

The cause cannot lie in the teaching of Classics to the boys 
of the Public Schools, for these form but a limited class, and 
for a good many years have enjoyed access to science teaching. 
I suggest that the causes rather lie in the strong reaction 
provoked by the extraordinarily one-sided results produced 
in the specially aided schools, which devoted themselves to 
Science in the last years of the reign of Victoria, and to the 
narrow formalism of much of the teaching. It was discovered 
that man cannot live and grow on science alone, and the 
revolt perhaps went too far, until the Great War forcibly 
reminded the nation that things were not well. 

It ought to be possible to think out a general education 
in which all will be able to gain some elementary insight 
into the workings of the physical universe and to come to 
understand the meaning of scientific method, and the point 
of view of the scientist. So much is necessary if in the future 


we are all to understand one another. It ought to be possible 
also so to teach Science that those who learp it do not become 
intolerant and unsympathetic, but realize that, though it 
is a necessary part, it is only a part of the modern citizen's 
outfit for life. 

This book is written by one who has realized this, and who 
knows how to teach with breadth and without exclusiveness. 
Its pages give information and provoke curiosity ; at many 
points they suggest that there are other realms of knowledge 
of a quite different sort. This characteristic, which may 
offend the purist, seems to me all to the good, and I hope 
that the work may find its way into the hands of many. 



If Science is to retain the honourable place it has won in 
the educational system of this country, I believe we shall 
have to recognize that it is the greatest of the " humanities," 
and deliberately abandon the so-called " utilitarian " stand- 
point. There are signs that this fact is being realized, and 
that schoolmasters are becoming alive to the vital truth 
recently re-expressed by Dr. Singer, " Science is a method 
and not a collection of facts." 

The present book is an attempt to present to students 
of the School and Higher Certificate standards a logical 
course of chemistry which shall acquaint them with modern 
ideas and give them an insight into the problems, methods 
and achievements of the science. I have emphasized the 
physical point of view as a ready means of building a frame- 
work, but I have, I hope, not sacrificed the useful (and enter- 
taining) parts of the usual school inorganic chemistry in doing 
so. I have not hesitated to mention uncommon substances 
where they are likely to thrill the youthful imagination, 
and I have included many biographical and historical facts, 
the psychological value of which as an aid in fixing the atten- 
tion, assisting the memory and arousing the enthusiasm of 
the student is well known to every teacher. 

It is my pleasant duty to record my heartiest thanks for 
much help received in the preparation of this book. I am 
especially indebted to Dr. H. J. H. Fenton, F.R.S., who 
generously undertook to read the whole of the proofs and 
whose suggestions have been of the greatest value. If I have 
caught a little of the spirit of his own Outlines of Chemistry 
my aim will have been amply achieved. 



My thanks are also due to Dr. W. A. Knight, of Marlborough 
College, who read the proofs and gave me much invaluable 
advice ; to Professor Soddy, F.R.S., for reacting and criticizing 
the chapter on the Structure of the Atom ; to the Master 
of Marlborough for writing the foreword ; to A. C. K. Toms, 
Esq., M.A., of Clifton College, who read the proofs and whose 
trenchant criticism was of much assistance ; and to some 
of my senior students, B. E. Berry, D. M. Stern, C. P. Wright 
and S. R. M. Porter, for preparing and checking the calcu- 

I have also to thank Professor Sir Ernegt Rutherford, 
F.R.S., for allowing me to quote from his Bakerian Lecture, 
1920 ; Messrs. G. Bell & Sons, Ltd., for the loan of blocks ; and 
Messrs. Longmans, Green & Co., Messrs. J. & A. Churchill, and 
Mr. John Murray, for kind permission to quote from Stewart's 
Recent Advances in Inorganic Chemistry, Moliriari's Chemistry, 
and Soddy's Interpretation of Radium respectively. 

In conclusion, perhaps I may be allowed to say that this 
is the third generation of Clifton chemical textbooks pub- 
lished by Messrs. Edward Arnold & Co. I hope that the 
connection may continue to be successful ! 




In the nine years that have elapsed since the first edition 
of this book was originally published, many advances have 
been made in both pure and applied chemistry. Such of these 
advances as lie within the present limits have been in- 
corporated, but care has been taken to preserve the character 
and individuality of the book with as little change as was 
consistent with the necessary additions. In particular, the 
temptation to include much more theoretical matter has been 
withstood, and it is hoped that readers will recognize in the 
new edition not a new book recast on a different basis, but 
an up-to-date form of the original edition. 

The author's warmest thanks are due to those very numerous 
correspondents who, since 1922, have written to him pointing 
out difficulties or inaccuracies and making suggestions for 
improvements. To all of them he believes that he has replied 
personally, but he would now take the opportunity of renewing 
his thanks. Not the least of the benefits that the book haa 
brought him is a large number of personal friends. 

The author desires to express his thanks to Mr. R. B. 
Pilcher for permission to reproduce the portrait of Gay- 
Lussac and to Mr. 1\ W. Clifford of the Chemical Society 
for those of Faraday and Berzelius. 






Headmaster of Harrow School . * . . v 





I HISTORICAL ......*! 























































CHROMIUM . . . . .483 

IODINE . . . . .491 



INDEX OF NAMES . . . . . .563 

SUBJECT INDEX ...... 568 

ANSWERS . . . . . . .574 

WEIGHTS . , 576 






III ROBEBT BOYLE ....... 27 

JV JOSEPH Louis GAY-LUSSAC . . . . .37 










Many suggestions have been made as to the origin of the 
word Chemistry, but that which has most to support it derives 
the word from the ancient name for Egypt, khem or chem* 
which means " black," and was given to Egypt on account of 
the dark soil of that country. Whether this be the true 
derivation or not, it is at least certain that the ancient Egyp- 
tians were acquainted with more chemical operations than any 
other nation of antiquity, and that therefore the " Egyptian 
Science " as a name for chemistry is very appropriate. Egypt, 
however, had no monopoly of chemistry, for the Chaldeans, 
Chinese and Hindoos had acquired much chemical knowledge 
of an empirical nature, and it seems very probable that 
each nation of the ancient world contributed its share to the 
development of the science. 

The kind of chemical knowledge possessed by men in those 
remote times is very much what we should expect, although 
we may be inclined to underestimate it. Metallurgy, glass- 
making, dyeing, the manufacture of pigments and poisons, 
soap-making and embalming, together with the preparation 
of drugs for medicinal purposes, were the principal subjects 
that engaged the attention of the ancient chemists. 
Chemistry was, therefore, mainly practical and empirical ; 

1 B 


theory lagged behind and was neither closely connected with 
practice nor supported by experiment. 

The first theoretical chemistry of importance so far as our 
records go was that of the Greek chemists. PYTHAGORAS, 
DIOSCORIDES and ZOSIMUS, as well as many more, covering 
a period of about 1,000 years from 600 B.C., all made chemical 
speculations and occasionally practical chemical observations 
as well. It appears, however, that the empirical chemical 
facts amassed by the Greeks were largely obtained by obscure 
men whose names have not come down to us. It might, 
indeed, be very naturally expected that they were craftsmen 
metal-workers, dyers, and the like and it is well known 
that the educated Greek had a marked distaste for experi- 
mental work, which did not appeal to his national genius. 

The Greek theories of the structure of matter made the 
transmutation of the metals seem to be quite possible, and 
observations of the deposition of copper on an iron blade 
placed in a solution of vitriol, and many similar effects, must 
almost inevitably have been interpreted as actual examples 
of transmutation. If transmutation was possible, it was 
obviously a fact of supreme importance, for gold then, as 
now, was the sinews of war and the passport to ease and 
luxury. It is not surprising, therefore, to find that the 
problem of converting " baser " metals into gold soon began 
to occupy a prominent place in chemistry and finally engrossed 
so much of the attention of chemists that " chemistry " 
became practically synonymous with " gold-making." 

When the empires of Byzantium and Persia were over- 
thrown by the armies of Islam (seventh century A.D.), the 
Muslim conquerors, after they had settled down, began to 
encourage learning, and to that end had translations made 
of all such important Greek works as were available to them. 
These translations were made either directly from Greek into 
Arabic, or more often at least at first from Greek into 
Syriac and then into Arabic. Greek chemistry was thus 
introduced to the Arabs, who soon developed a passion for 
the science, and as they were also masters of Egypt they 


>^5 ^JA$\,$V\^\M\ W^\^jfi>&& 

\ ^ ( ^y^u\^'.'^>^>r^^^ 

lb: ~~\ -J "- ^- --- ,-__.... .^s 

Fio. 1. Page of Arabic Chemical MS., dealing with distillation. 



were able to effect the necessary union of the theory of Greece 
with the practice of Egypt. This in itself was a great advance. 
The art of experiment was practically unknown in Greece, 
and the correlation of it with theory and therefore the 
establishment of scientific method is the great achievement 
of the chemists of Islam. 

Traces of Arabic influence on chemistry are still to be found 
in many of our chemical terms alchemy, for example, is 
merely " chemistry " with the prefix a/-, the Arabic word 
for " the ", while alembic, a kind of retort, aludel, a pecu- 
liarly shaped receiver, and alcohol, from the Arabic al-khul, 
all remind us of the days when the Muslims were the first 
chemists of the world. Al-kuhl was the name given to the 
black powder with which the Muslim ladies blackened their 
eyebrows and eyelids possibly lead and antimony sulphides : 
its transference to spirit of wine is a strange fact for which 
PABACELSUS is responsible. 

The greatest of the Arabic chemists were JABIB IBN HAYYAN, 
whose name has been westernized to GEBEK, and who lived 
in the eighth century A.D., ABU BAKB MUHAMMAD IBN ZAK- 
ABIYYA AL-RAZI (died A.D. 925), ABU'L-HASAN f ALi AL- 
ANDALUSI, sometimes known as IBN ABFA C RA'S (died A.D. 
teenth century), and c lzz AD-DiN AIDAMIB IBN e Au AL- 
JILDAKI (died A.D. 1361 or thereabout). Of these, Jabir has 
the greatest reputation, but AL-RAZI (known to the West as 
Rhazes) was perhaps a better practical chemist. Jabir is 
said to have written some 500 treatises on chemistry, of which 
only a few are now known to exist. Ibn Arfa c Ra's wrote a 
long chemical poem, known as Particles of Gold, of which 
many copies are preserved in our museums, while Abu'l- 
Qasim al- c lraqi wrote the important work Al-Muktasab, or, 
to give it its full title in English, Knowledge Acquired Concern- 
ing the Production of Gold, in which we find a very full account 
of the chemical theory of the day and the experimental work 
that the author had carried out in connection therewith 

The westernmost province of the Muslin Empire was Spain, 
and it was here that chemistry first took firm root in Europe. 


Fostered by the Muslim rulers, science in Spain flourished, 
and chemistry might at the present day have been in a much 
higher stage of development if the great disaster of the defeat 
of the Muslims by CHABLES THE HAMMER at Tours in A.D. 732 
had not taken place. This triumph of the forces of reaction 
was a great blow to science in general and chemistry in par- 
ticular, and although the Moorish power in Spain was not 
finally overthrown until 1492, irreparable damage had been 

From Spain, the study of chemistry gradually spread 
over the rest of Europe, and the European chemists of the 
Middle Ages frankly admitted their debt to the Muslims. 
They did not, however, realize what they owed to the Greeks 
until after the Renaissance, when Greek manuscripts were 
first directly available. It was then possible to see the 
modifications of Greek chemistry produced by its passage 
through Islam. 

When chemistry became thoroughly established in Europe 
rapid advance took place, largely owing to the more systematic 
mind of the European as compared with that of the Asiatic. 
The faculty of correlation is not highly developed among 
Eastern nations, and this faculty is of course a very essential 
one for scientific progress. We find that the Latin works 
of European origin of the thirteenth and fourteenth centuries 
show a marked advance in systematic arrangement, and this, 
with the scientific method developed by the Muslims, soon 
bore fruit. 

The great chemists of the Middle Ages in Europe were 
ROGER BACON (1214-1292) and PARACELSUS (1493-1541). 
By the time of Paracelsus, the obsession of metallic transmuta- 
tion was on the wane, but the change was not altogether one 
for the better, for chemistry became temporarily the handmaid 
of medicine, and the efforts of chemists were mainly directed 
to the preparation and investigation of drugs. This was the 
period of latrochemistry, or medical chemistry, arid may be 
taken as covering roughly the years A.D. 1500-1700. 

Towards the end of this time appeared the famous ROBERI 
BOY.LB (1627-1691), who lived in the reign of Charles II and 


was one of the founders of the Royal Society. Boyle was a 
scientist of the first rank, and his book The Sceptical Chymiat 
(1661) was an extremely valuable contribution to chemical 
theory. In it he insists upon a rigid observance of scientific 
method, experiment, observation and inference ; he questions 
the truth of the old theory of the constitution of matter, which 
regarded Fire, Air, Earth and Water as the four " elements," 
and defines an element as a substance that cannot be split 
up into other substances ; and finally he suggests that matter 
is composed of small particles of different shapes and sizes 
called atoms, combination and separation of which take place 
in chemical changes. 

Boyle was, however, in advance of his age, and his ideas 
had not the immediate great effect which might have been 
expected. Chemists were very much engrossed at this time, 
and for a hundred years after, in the study and expansion of 
a theory of combustion. JOHANN JOACHIM BECHER (1635- 
1682) suggested in his Physica Subterranea (1669) that the 
process of combustion was a decomposition of the burning 
substance into its constituents. All substances that would 
burn he assumed to contain a fatty substance, the terra 
pinguis, and supposed that on combustion this substance 
was lost. The residue after combustion represented the 
remaining constituent or constituents of the original body. 
GEORG ERNST STAHL (1660-1734) developed this theory and 
named the substance that was lost on combustion phlogiston, 
the " principle of fire," a word that had been previously 
used by Boyle but not in so clear and definite a way. The 
Phlogiston Theory offered a simple and coherent explanation 
of many facts of combustion, and we shall return to it later 
(p. 410). It will be sufficient here to say that the Phlogiston 
Theory may well be regarded as the first theory of Modern 
Chemistry, and greatly helped the advance of the science. 

Towards the end of the eighteenth century the Theory of 
Phlogiston was overthrown, to some extent by the work of 
PRIESTLEY (1733-1804) (quite unintentionally on his part, for 
he was a confirmed phlogistian), but chiefly by the work of 
ANTOINE LAURENT LAVOISIER (1743-1794). It had for long 


been known that the residue left after the combustion of a 
metal weighed more than the metal taken originally, but 
the knowledge was not general arid its importance not per- 
ceived. Lavoisier, however, after many years of patient and 
brilliant experimental work, was able to convince the chemical 
world that, probably in all cases of combustion and certainly 
in most, an increase of weight occurred, whereas according 
to the Theory of Phlogiston a decrease would be expected. 
These observations made the Phlogiston Theory totter, and it 
fell when Lavoisier was able to show that Priestley's newly 
discovered gas, dephlogisticated air, was present in the atmo- 
sphere and, in all cases of combustion in air, combined with 
the burning substance, whence the increase in weight. He 
later called this gas oxygen. Lavoisier was guillotined during 
the French Revolution, in 1794. 

A few years later JOHN DALTON (1766-1844) made the 
greatest step in the history of chemistry by his work 
on the Atomic Theor}^. We shall consider this in some detail 
in the next chapter and therefore need say no more here 
than that with the aid of this theory and the new theory of 
combustion just perfected by Lavoisier, the nineteenth cen- 
tury opened for chemistry under the happiest auspices. Dur- 
ing the next 130 years an advance was made that can modestly 
be described as marvellous. 

The history of chemistry since 1800 cannot be adequately 
treated in the present book, and has already been excellently 
described by many distinguished chemists to whose books 
(some of which are mentioned at the end of the chapter) the 
reader is referred for further information. The most note- 
worthy featxires, however, are the development of organic 
chemistry by LIEBIG, KEKULE, EMIL FISCHER and many 
others of scarcely less celebrity ; the rise of stereochemistry, 
due largely to the labours of PASTEUR, VAN'T HOFF and 
SIR WILLIAM POPE ; the rapid progress of physical chemistry, 
in which the names of RAOULT, ARRHENIUS, OSTWALD and 
NERNST are perhaps the most prominent ; the discovery of 
the rare gases of the atmosphere by SIR WILLIAM RAMSAY ; 
the discovery of radioactivity by BECQUEREL and M. and 


MMB. CCTRIB ; and the recent work on the structure of the 
atom and the nature of valency by SIB J. J. THOMSON, LORD 
DIRAC, ASTON and many others. 


E. von Meyer, History of Chemistry. 
J. Campbell Brown, History of Chemistry. 
T. M. Lowry, Historical Introduction to Chemistry. 
M. Berthelot, La Chimie au Moyen Age. 

E. von Lippmann, Die Entstehung und Ausbreitung der Alchemie. 
Alembic Club Reprints. 

See also E. J. Holmyard, Chemistry to Dalton, 1925. The Great Chemists, 
1928. Makers of Chemistry, 1931. 


1. Write an account of the early history of chemistry. 

2. Who were the chief Muhammadan chemists ? What contribu- 
tions to the development of chemistry were made by Islam ? 

3. Mention the chief landmarks in the progress of chemistry since 
the time of Boyle. 

4. With what discoveries do you associate the names of (a) Becher ; 
(6) Boyle ; (c) Priestley ; (d) Lavoisier ; (e) Dalton ? 

6. Estimate the probable advance of chemistry in the next fifty 


The idea that matter is composed of extremely minute 
particles, the various kinds and arrangements of which give 
rise to the many forms that matter assumes, was first 
elaborated by the Greek philosophers, an excellent account 
of whose work is to be found in The Study of Chemical Com- 
position by the late Miss FREUND. In the ancient world, 
however, theories of the constitution of matter could be 
nothing more than speculations or, to put it plainly, guessea 
entirely unsupported by experimental, if not by observa- 
tional, evidence. We do not call MOTHER SHIPTON the 
inventor of the motor-car, although she prophesied that some 
day carriages would move without horses, and in just the 
same way we do not regard the Greek philosophers as the 
founders of the Atomic Theory. That position is occupied 
by JOHN DALTON (1766-1844). 

Dalton was a Quaker schoolmaster, born at Eaglesfield in 
Cumberland. Like all other scholars of the time he was given 
a classical education (although he studied mathematics as 
well), and was probably therefore acquainted with the classical 
ideas on the structure of matter. He seems to have pondered 
over these, and the more definite ideas of the atomic nature 
of matter suggested by NEWTON, and finally reshaped the 
theory in such a form that it explained certain facts already 
known and was also capable of experimental proof or 
disproof, partial if not complete. It is important to note 
that Dalton did not arrive at the Atomic Theory from a 
consideration of experimental results already gained ; he 

9 B* 


thought out the theory first and then made experiments 
to test it. 

The main points of the Atomic Theory *as formulated by 
Dalton are as follow : 

1. Matter is composed of a great number of extremely small 
particles, called atoms. [The word atom means " indivisible," 
and was given to these particles because they were considered 
to be incapable of further division, in the case of elements. 
The " compound atoms " of compounds on division, of course, 
would give atoms of elements, and therefore in this sense a 
compound atom is also indivisible, or is at least the smallest 
particle of the compound of which one can conceive.] 

2. All the atoms of the same element are identical in all 
respects, and different from those of other elements. 

3. Atoms are indestructible and cannot be created. 

4. In the formation of compounds, combination occurs 
between small whole numbers of atoms of the elements 

5. All the " compound atoms " of a compound are exactly 

From these postulates (they were at that time nothing 
more) it is possible to deduce certain conclusions that can 
be tested by experiment. This at once put the Atomic 
Theory on a sound scientific basis, and explains the fact 
that chemists universally consider Dalton to be the Founder 
of the Theory. No scientific theory is of any use unless 
it can be tested experimentally, either directly or indirectly. 
In spite of this fact, which should be sufficiently obvious, 
we shall find that many " untestable " theories have been 

The first deduction to be made from the Atomic Theory 
is called the Law of Constant Composition, viz., that the 
same chemical compound always consists of the same elements 
combined together in the same proportion by weight. This 
follows from 2, 4, and 5 above. As this is the first " Law " 
of chemistry we have met with, it will be well for us to con- 
eider exactly what we mean by the word law in this con- 
nection. It has clearly a quite different meaning from that 

By permission of the Manchester Literary and Philosophical Society. 



implied in the phrase " penal law," for it would be ridiculous to 
suppose that up to the time of Dalton chemical compounds 
led a free and easy existence, and that water, for example, 
could consist of tin and lead, if it so desired, at one time, and 
of hydrogen and oxygen at another. It would be equally 
foolish to imagine that if a chemical compound had the 
temerity to break the " law " of constant composition it could 
be punished in any way (although many chemists of the 
middle of the nineteenth century must have heartily wished 
they could sentence nitric oxide to a term of hard labour for 
its effrontery in refusing to obey the Law of Even Numbers). 
We have discovered, therefore, what a chemical " law " is 
not, but we still have to decide what it is. We can easily 
do this if we take a simpler example. It is a " law " of 
Nature that every man has two eyes. Now this does not 
prevent a man from plucking out his eye if it offend him, 
neither does it follow that a man may not have a third (glass) 
eye in his waistcoat pocket, and indeed he has a third rudi- 
mentary eye in the top of his head, so zoologists tell us. It 
does, however, imply that, as a matter of fact and observation, 
men do have two eyes. It is a short way of saying " A has 
two eyes, B has two eyes, C, D, E, and 1,900,000,000 more 
human beings have each two eyes." It is, in fact, a summary 
of experience, and this is true of chemical laws. The Law of 
Constant Composition means simply this : that up to the 
present no compound has been discovered of a variable composi- 
tion. If such compounds were ever discovered, the " law " 
would vanish, and here again we see the difference between 
a penal law and a chemical law, for if I am fined 1 for 
allowing a dog to stray, I find that after the infraction the 
law against the straying of dogs is quite unharmed. 

Let us now examine the relation between the Law of 
Constant Composition and Dalton's Atomic Theory. In 
essence it is this : if the Atomic Theory be true, then, as a 
matter of fact and observation, the composition of any one 
chemical compound should always be the same. How far is 
the converse true ? Does it follow that if, by experiment, 
we find that the composition of a certain chemical compound 


is constant, therefore the Atomic Theory is true ? A little 
reflection will convince us that this is not so. Suppose a man 
to be going by train ; well, we can form a Travel Theory 
about him. We should, as a matter of fact, expect him to 
come along the road to the station, to buy a ticket, to arrive 
at the station before the scheduled time of departure of the 
train, to hurry if he were late, and possibly to carry a bag. 
Now, suppose we see a man walking towards the station. 
Does it follow that he is going by train ? Of course it does 
not ; he may be going to meet a friend, or to claim damages 
for a broken article, or for many other reasons. But the 
fact that he is going towards the station does agree with the 
theory that he is going by train ; it by no means proves it. 
Suppose, however, that he is running as well, that it is nearly 
time for a train to go, that he is carrying a bag, and that he 
goes to the booking-office and buys a ticket is the theory 
proved now ? No, for he may be acting for the cinemato- 
graph or he may be mad ; we cannot be absolutely certain 
that he is going by train until we have seen him get into it 
and go off in it. But the above evidence would make it 
extremely probable, and the Travel Theory would be much 
more reasonable than the Cinematograph or Insanity Theories, 
and we should be justified in accepting it in the absence of 
definite evidence for either of the others. 

This is exactly the position that the Atomic Theory was 
in all through the nineteenth century. No one, so to speak, 
had seen the man get into the train and go off, but the cir- 
cumstantial evidence was very strong indeed. There were no 
" eye-witnesses " until the early years of the present century. 

The fact that, in practice, the Law of Constant Com- 
position is found to be true, is then in accordance with the 
truth of the Atomic Theory, but it does not prove it. It is 
circumstantial evidence. Now the value of circumstantial 
evidence increases, as LORD DARLING has pointed out, not 
in arithmetical but in geometrical progression, so that if 
we find the next deduction from the Theory to be borne out 
in practice, the probability of the truth of the Theory will be 
greatly increased ; and if several deductions are tested and 


found satisfactory the probability will become a practical 

Before we pass on to the next deduction from the Atomic 
Theory there is another point to be considered. One is 
sometimes asked how to prove the truth of the Law of Con- 
stant Composition, or of other chemical laws. The answer is 
that the question is badly expressed. It is clearly impossible 
to prove a law of chemistry in the general sense, for to " prove " 
the Law of Constant Composition, for example, we should 
have to analyse all the specimens of all the compounds that 
exist now, have ever existed, and will exist in the future in 
the whole universe. The most we can do is to show that it 
is true in certain cases, and to tell the sceptic that if he knows 
of an exception, to produce it. 

Incidentally, it has to be remarked that all measurement is 
bound to be more or less inexact, and that we can therefore, 
even in a special case, demonstrate the truth of the Law 
only within the limits of experimental error. 

The Law of Multiple Proportions is the next conclusion 
we can draw from the Atomic Theory. It is that when two 
elements combine to form more tJian one compound, then the 
weights of one of those elements which combine with a constant 
weight of the other are in a simple ratio to one another, and is 
a logical inference from 2 and 4 (" small " whole numbers). 
Suppose, for example, that the elements A and B unite 
together to form two different compounds. The simplest 
imaginable case will be when in one of the compounds 1 atom 
of A combines with 1 atom of B, and in the other compound 
1 atom of A combines with 2 atoms of B. Since, in one 
" compound atom " of each of the two compounds there is 
1 atom of A, it follows, from Dalton's assumption 2, that the 
weight of A is constant in these two " compound atoms/' 
The weights of B, on the other hand, will be in the ratio of 
the number of atoms of B present in each " compound 
atom " ; m this case, 1:2. This is a simple ratio, and if 
compounds are always composed of small numbers of atoms, 
the ratio always must be simple. It does not matter of which 
element we take the constant weight, nor what weight we 


take as the constant weight, as long as we find the ratio to 
one another of the weights of the other element that have com- 
bined with this constant weight. Thus, suppose in the above 
instance we take the weights of A that have combined with 
a constant weight of B. Let 1 atom- weight of B be the 
constant weight of B. Then in the first compound we have 
1 atom-weight of A combining with the given, constant, weight 
of B. In the second compound, 2 atom-weights of B combine 
with 1 atom-weight of A, therefore 1 atom-weight of B (the 
" constant " weight, or " reference " weight) would combine 
with J an atom-weight of A. Now the ratio of 1 atom-weight 
of A to ^ atom- weight of A is 2 : 1. 

Having deduced the Law, Dalton's next care was to test 
it experimentally. He worked on certain compounds of 
carbon and hydrogen and found that he obtained the expected 
results. It is, however, easier to illustrate the truth of the 
Law with other compounds, such as the oxides of lead or 
copper. To make quite certain of understanding the Law it 
is best to work through several numerical examples. 

Examples. (1) Phosphorus trichloride and phosphorus ponta- 
chloride contain 77-45 per cent, and 85-13 per cent, chlorine respec- 
tively. Show that these compounds illustrate the law of multiple 

In the trichloride 

77-45 gms. chlorine are combined with 22-55 gms. phosphorus. 

In the pentachloride 

85-13 gms. chlorine are combined with 14-87 grns. phosphorus. 

therefore - X 22-55 gins. ,, M ,, 22-55 gms. ,, 


129-1 gms. ,, ,, ,, 22-55 gms. ,, 

Thus in those two compounds the weights of chlorine that are 
combined with a constant weight (22-55 grams) of phosphorus are 
77-45 gms. and 129-1 gms. respectively ; these weights are in the ratio 

(2) Four oxides of titanium contain the following percentages of 
oxygen : 

Monoxide . 24-96 per cent. Sesquioxide . 33-29 per cent. 
Dioxide . . 39-95 ,, Trioxide . . 49-9* 

Show that these figures illustrate the law of multiple proportions. 


In the monoxide 

75-04 gms. titanium are combined with 24-96 gms. oxygen. 
In the sesquoixide % 

66-71 gms. titanium are combined with 33-29 gms. oxygen 

// P7 

therefore - x 24-96 gms. 24-96 gms. 


50-02 gms. 24-96 gms. 

In the dioxide 

60-05 gms. titanium are combined with 39-95 gms. oxygen. 

therefore |2^5 X 24-96 gms. 24-96 gms. 

5y "i'o 

37-52 gms. 24-96 gms. 

In the trioxide 

50-05 gms. titanium are combined with 49-95 gms. oxygen. 

therefore '^5 x 24-96 gms. 24-96 gms. 

25-01 gms. 24-96 gms. 

Thus the weights of titanium combined with a constant weight (24-96 
gms.) of oxygen are 75-04 gms., 50-02 gms., 37-52 gms., and 25-01 gms. 
respectively ; these are in the ratio 6:4:3:2. 

The Law of Reciprocal Proportions also follows as a 
corollary of Dalton's Atomic Theory. It states that if an 
element A combines with an element B and also combines with 
an element 0, then if B and C combine together, the proportion 
by weight in which they do so will be simply related to the ratio of 
the weights of B and C which (separately) combine with a constant 
weight of A. For example, sodium will combine with chlorine 
in the proportion of 1 gram of sodium to 1-54 grams of chlorine. 
Sodium will also combine with iodine in the proportion of 
1 gram of sodium to 5-52 grams of iodine. Therefore if 
iodine and chlorine combine together, the ratio of the weights 

in which they do so should be simply related to y^p-r, which 

is the ratio of the weights of these elements that combine 
with a constant weight of sodium. Now as a matter of 
experimental fact, it is found that iodine and chlorine will 
combine together, in the proportion of 1 gram of chlorine to 

1*19 grams of iodine. IB the ratio c rj simply related to 



1*19 5*52 

-y- ? If we simplify the ratio by dividing top and 


bottom by 1-54 we get , when it becomes apparent that 

3-58 . 1-19 n . . , 

j is to : as 3 is to 1 a simple ratio. 

The logical necessity of this law will be obvious when it is 
remembered that Dalton postulated that (a) combination 
takes place between small whole numbers of atoms, to form 
a compound, and that (6) all the atoms of the same element 
are of exactly the same weight. Note that the Law does not 
state that the elements " B " and " C " will combine ; it 
states only what will happen if they do. 

A final deduction from the Atomic Theory is the Law of 
the Conservation of Matter, which states that matter is 
uncreatable and indestructible. This Law is clearly essential 
and fundamental for the development of chemistry as an 
exact science. Logical proof of the Law cannot, of course, 
be attempted ; the most one can do is to show that in any 
special case of chemical change, the total weight of the 
products is exactly equal (within the limits of experi- 
mental error) to the total weight of the substances started 

Careful experiments of this kind were carried out by 
LANDOLT (1909) and MANLEY (1912). Both were unable to 
detect the slightest change in weight during the chemical 
reactions they investigated. 

We shall find later on (Chapter XVI) that Dalton 's Theory 
has proved to be inexact in every particular, but that as a 
working guide for all but the most minutely accurate research 
it is not likely to be superseded. We have only to think of 
the wonderful development of chemistry during the nineteenth 
century to realize how useful the theory has been. 


1. What are the main features of the Atomic Theory as formulated 
by Dalton ? 

2. What do you understand by a law in chemistry T 


3. State the Law of Multiple Proportions and say how you would 
illustrate its truth. 

4. Show that the laws of constant composition, reciprocal propor- 
tions and multiple proportions are direct corollaries of the Atomic 

5. The three oxides of phosphorus contain the following percentages 
by weight of phosphorus : 

Peritoxide . . . . . .43 668 per cent. 

Tetroxide ...... 49-212 ,, 

Trioxide 56-365 

Show that these figures illustrate the Law of Multiple Proportions. 

6. 10 gms. stannous sulphide contain 7-877 gms. tin, while 8 gms. 
tannic sulphide contain 5-196 gms. tin. 

Find the weight of sulphur combined with 1 gm. tin in each case and 
show that your results agree with the Law of Multiple Proportions. 

7. Samples of the three chlorides of vanadium contained the following 
weights of the elements : 

Vanadium. Chlorine. 

(i) 0-3643 gm. 0-5060 gm. 

(ii) 1-6149 grn. 3-3638 gms. 

(111) 0-6580 gm. 1-8279 gm. 

Calculate the weight of chlorine combined with 1 gram of vanadium 
in each case and show that the results illustrate the Law of Multiple 

8. Ferrous sulphate contains 36-76 per cent, iron and ferric sulphate 
contains 27-93 per cent. iron. Show from these figures that the law 
of Multiple Proportions is obeyed in the case of the sulphates of iron. 

9. Samples of four oxides of lead contain the following weights of 
the elements : 

Lead. Oxygen. 

Suboxide ..... 4-144 gms. 0-160 gm. 

Monoxide 1-036 gm. 0-080 gm. 

Sosquioxide 5-2007 gms. 0-600 gm. 

Peroxide 6-488 gms. 1-002 gm. 

Show that those results illustrate the Law of Multiple Proportions. 

10. 13-962 gms. of aurous chloride were taken and acted upon with 
water. This gave 7-888 gms. of gold together with some auric chloride. 
The auric chloride was then heated to 220 and the resulting metal 
was found to weigh 3-944 gms. 

Find the weight of chlorine combined with 1 gram of gold in each 
of the chlorides and show that your answer illustrates the Law of 
Multiple Proportions. 

11. Given a sample of iron wire, what experiments would you make 
to find whether iron oxides obey the Law of Multiple Proportions ? 



" Substances react only by definite weights." AL-JILDAK! 
(| ca. 1360 A.D.) 

If we take a definite weight of pure sulphuric acid, say 
10 grams, and dilute it with water, we shall find that it will 
act upon zinc. The zinc will dissolve arid an inflammable 
gas, hydrogen, will be evolved. We shall also find that the 
weight of zinc that 10 grams of sulphuric acid will dissolve 
is limited and not indefinite. Experiment shows that 10 
grams of sulphuric acid, when diluted with water, will dissolve 
6*63 grams of zinc ; if more zinc be added the excess will 
remain undissolved, and if less be used, some of the sulphuric 
acid will be left over. We therefore say that 10 grams of 
sulphuric acid are equivalent to 6-63 grams of zinc. The 
hydrogen, too, that is evolved when 6-63 grains of zinc are 
dissolved in 10 grams of sulphuric acid, diluted with water, 
is found always to weigh 0-204 grams ; we may therefore 
say that 0-204 grams of hydrogen are equivalent to 10 grams 
of sulphuric acid or to 6-63 grams of zinc. 

The same sort of thing is true for all other reactions, and 
this constancy of reacting weights of elements and compounds 
is one of the fundamental facts of chemistry. It was men- 
tioned by AL-JILDAKI about 1360, and was re-discovered 
by CAVENDISH in 1766, who showed that the weights 
of sulphuric acid and nitric acid that would separately 
neutralize a definite weight of caustic potash, would also 
separately dissolve a (different) definite weight of marble. 
The weights of sulphuric acid and nitric acid he therefore 
called equivalent to one another ; the weights of potash and 



marble were also equivalent to one another and to the weights 
of the respective acids. 

Chemists soon found it convenient to choose a standard 
weight of a suitable substance in terms of which the equivalent 
weights of all other elements and compounds could be 
expressed. For this purpose they fixed upon one unit 
weight of hydrogen. On this system of reference, the equiva- 
lent of sulphuric acid is 49, since when hydrogen is liberated 
from sulphuric acid, 49 grams of the acid always yield 1 gram 
of hydrogen ; or when 49 ounces of the acid are used, they 
yield 1 ounce of hydrogen, and so on. We can therefore 
define the Equivalent of any element or compound as that 
number of units of weight of it which will react either directly 
or indirectly with one of the same units of weight of hydrogen. 
Note that the equivalent is a number only ; it is incorrect to 
Bay that the equivalent of sulphuric acid is 49 grams the 
equivalent is not 49 grams but just 49. When we say 
that the equivalent of sulphuric acid, then, is 49, we mean 
that 49 grams, grains, ounces, pounds or tons of it will react 
with evolution of 1 gram, grain, ounce, pound or ton respect- 
ively of hydrogen. Similarly, the equivalent of oxygen is 
the number of units of weight of it which will combine with 
one of the same units of weight of hydrogen. Thus, although 
the unit of weight commonly used in chemistry is the gram, 
the equivalent of a substance is a number only, and, in the 
words used by AL-'IRAQI when speaking of another matter, 
" Know this therefore, and ponder thereon, for verily it is one 
of the fundamental things of this Divine Science." 

When we come to consider the equivalent of a metal, 
we could determine it by finding the number of grams of 
the metal that will liberate 1 gram of hydrogen from a 
dilute acid ; but there are many metals that do not dissolve 
in acids with liberation of hydrogen. In these cases we 
attack the problem by an indirect method. Oxygen will 
combine directly with hydrogen and we can therefore easily 
determine its equivalent directly. Experiments have shown 
that 8 grams of oxygen will combine with 1 gram of hydrogen ; 
hence the equivalent of oxygen is 8. Now, although it may 


be impossible to get the metal to liberate hydrogen from a 
dilute acid, it is almost always possible to make it combine 
with oxygen, and the number of grams of it which will com- 
bine with 8 grams of oxygen is taken as its equivalent, since 
8 grams of oxygen combine with 1 of hydrogen. In short, 
we assume that the weights of two substances which are equiva- 
lent to a constant weight of a third are also equivalent to one 
another, and this assumption has justified itself in practice. 

Note that the term equivalent weight is not confined to 
elements ; equivalents of compounds may also be determined 
and are included in the same definition. 
Methods of Determining the Equivalents of Elements. 
A. Metals. 

(i) If the metal is soluble in a dilute acid (or in a solution 
of an alkali) with liberation of hydrogen, take a known weight 
of it, dissolve it in a suitable dilute acid (or alkaline solution) 
and find the weight of hydrogen evolved. 

(ii) If the metal is not soluble in a dilute acid, convert a 
weighed quantity of it into its oxide, and find the weight of 
oxygen taken up (= increase in weight). Then calculate 
the number of grams of it which would combine with 8 grams 
of oxygen. 

(iii) If neither of the above methods is suitable, convert 
a weighed quantity of the metal into its chloride. Weigh 
this and calculate the number of grains of the metal which 
would combine with 35-5 grams of chlorine, since 35-5 is the 
equivalent of chlorine. Methods ii and iii may also be used 
in the reverse way, that is, reduction of the oxide or chloride 
to the metal. 

(iv) The equivalent of a metal may often be conveniently 
found by displacement of it from a solution of one of its salts 
by means of a weighed quantity of another metal of known 
equivalent. Then the equivalent of the first metal is that 
number of grams of it which is displaced by the equivalent 
in grams of the second. 

(v) By electrolysis. It was shown by FARADAY that in 
electrolysis the weights of substances liberated at the elec- 
trodes were in the ratio of their chemical equivalents. By 


measuring, therefore, the weight of a metal deposited in an 
electrolytic cell and the weight of hydrogen liberated by the 
same quantity of electricity, the equivalent of the metal may 
be found. 

(vi) A compound of the metal with an element of known 
equivalent may be converted into another compound of 
the metal with other elements of known equivalents. From 
the weights of the two compounds the equivalent of the 
element may be determined. For example, calcium chloride, 
containing the metal calcium, whose equivalent is desired, 
and chlorine, equivalent 35-5, is converted into calcium 
sulphate, containing calcium, sulphur and oxygen. We 
know from experiments with sulphuric acid that the equiva- 
lent of the " sulphate " group is 48. Suppose we took 
1-000 grams of calcium chloride and obtained 1-225 grams of 
calcium sulphate. Then if x be the equivalent of calcium 
x + (35-5) 1-000 

/. l-225z + (35-5) (1-225) = x + (48) 

.'. 0-225s = 48 - 43488 = 4-512 

__ 4-512 
*' X ~~ (H225 
= 20-06. 
B. Non-Metals. 

The determination of the equivalents of non-metals, while 
equally simple in theory, is usually more difficult in practice, 
since so many of the non -metallic elements are gases that 
need more skill in manipulation. 

(i) By direct combination of the element with hydrogen, 
when the equivalent may be calculated from the weight of 
the element taken and the weight of the product. 
(ii) By formation of the oxide or chloride. 
(iii) By method vi as given for metals. 
Details of the determination of the equivalents of various 
non-metallic elements will be found in the descriptive part 
of this book. 

The exact determination of equivalents is of great import- 
ance, since upon it mainly depend the values of the atomic 


weights of the elements (Chapter VI). It is usual at the present 
day to take as standard not 1 unit of weight of hydrogen but 
8 units of weight of oxygen. If the hydrogen unit be taken, 
the equivalent of oxygen is not exactly 8, but 7-94, hence 


equivalents on the oxygen standard are^ ~ times the equiva- 

/ *Ji 

lents on the hydrogen standard. The oxygen standard is 
preferred, since the numbers it gives are more often or more 
nearly whole numbers than those on the hydrogen standard. 
There is also theoretical justification for the oxygen standard, 
but the matter is too advanced to be discussed here. 

Equivalents of Compounds. The equivalent of a com- 
pound is determined by making use of some reaction in 
which the compound takes part. Thus, the equivalent 
of sulphuric acid is the number of grams of the acid which 
will yield 1 gram of hydrogen on treatment with zinc. If 
the equivalent of sulphuric acid be known, then the weight 
in grams of caustic soda which the equivalent in grams 
of sulphuric acid will neutralize is the equivalent of caustic 
soda. By continuing this process, the equivalent of any 
compound may be discovered. It follows that, if two 
substances react together, they will do so in the propor- 
tion by weight of their equivalents, since their equivalents 
have been determined in this way. This property is made 
use of in volumetric analysis, where a normal solution of 
a substance is one that contains the equivalent in grams 
of the substance in one litre of solution. Hence, equal 
volumes of all normal solutions are exactly equivalent to one 
another. Thus 10 c.c. of N-hydrochloric acid will neutralize 
exactly 10 c.c. of N-caustic soda or N-caustic potash ; and 

25 c.c. of or half -normal hydrogen peroxide will de- 


colourize exactly 25 c.c. of ~ potassium permanganate, and 

BO on. It is as a rule possible to find the equivalent of a 
compound from an inspection of its formula (Chapter VII), 
but you should carefully avoid the mistake of imagining 
that the equivalent of a compound is always the sum of 


the equivalents of its constituent elements. It often is and 
often is not. 

It should be noted that the equivalent df a substance may 
sometimes have different values for different reactions or in 
different compounds, but these values are always simply 
related to one another. Thus the equivalent of copper in 
black copper oxide is 31*8, while in brown copper oxide it is 
63-5, the ratio of the two being 1:2. The explanation of this 
phenomenon we shall discuss later. 


1. Define equivalent. What methods are available for determining 
the equivalents of metals ? 

2. 0-461 grams of a metal on oxidation gave 0-503 grams of its oxide. 
What is the equivalent of the metal ? 

3. 1-342 grams of the chloride of a metal wore reduced in a current 
of hydrogen. The metal thus produced weighed 0-753 grams. Cal- 
culate its equivalent. 

4. On dissolving 0-36 grams of a metal in a dilute acid, 0-013 grams 
of hydrogen were evolved. Find the equivalent of the metal. 

5. What do you mean by a normal solution ? How many c.c. of 
N/10 hydrochloric acid would be required to neutralize 19-3 c.o. of 
N/3 caustic soda T 


Changes of temperature and pressure alter the volume of 
a gas, and the laws expressing the relation between tempera- 
ture, pressure and volume have been formulated by BOYLE 

In 1662 Boyle stated that the volume of a given mass of gas 
varies inversely as the pressure upon it, if the temperature is 
constant. This is known as Boyle's Law, and may be 
written shortly as 

V oc p if T is constant. 

Gas pressure is measured in terms of the length of a column 
of mercury which the pressure could support ; thus, " a 
pressure of 100 mm." means a pressure that would support 
a column of mercury 100 mm. high. For higher pressures, a 
larger unit, the " atmosphere" is used ; this is equal to the 
average pressure of the earth's atmosphere at sea-level, viz., 
760 mm. The pressure of 760 mm. is also called Normal or 
Standard Pressure. 

Assuming Boyle's Law, if we know the volume of a gas at 
one pressure, we can calculate what volume it would occupy 
at another, if the temperature remain constant. Suppose a 
gas to occupy 100 c.c. at 500 mm. pressure. What volume 
would it occupy at 760 mm. ? Since the pressure will be 
greater, the volume will be less, and therefore we must put 
the larger pressure at the bottom in the expression 



100 X 500 

New volume = - 

= 65-7 c.c. 

If you always ask yourself first whether the final volume 
will be greater or smaller, and arrange the factors of your 
expression accordingly, you will never go wrong. 

Charles' Law states that, provided the pressure is con- 

stant, a gas will expand -^ of its volume at C. for every rise 

in temperature of 1 C. and will contract -^ of its volume at 

C. for every fall in temperature of 1 C. If we have 1 c.c. 
of gas at C. the volume, v, at t C. will therefore be 

1 + -^_ c.c., or if we have x c.c. at C., at t C. the new 

4Ut O 

volume x' will be equal to ^(1 + 070) c - c - Suppose 

t = 273C., what will the volume of the gas be at this 
temperature ? For every fall of 1 in temperature, the gas 

contracts ^r^ of its volume at C., therefore at 273 it 

will have contracted -^- of its volume at 0, that is, it will 

have no volume at all. This is shown too by the formula, 

_ 273\ 
1 _[- j or x(0) = 0. If there- 

fore we could cool a gas to 273 C. it would vanish. As a 
matter of fact, all gases evade the issue by changing to liquids 
before this temperature, called absolute zero, is reached. The 
lowest temperature yet recorded is 0-75 above absolute zero, 
or 272-25 C. It can be shown theoretically that at absolute 
zero every substance would be perfectly devoid of heat ; 
hence a lower temperature is inconceivable. 

The Absolute Scale of Temperature uses the same degrees as 
the Centigrade Scale (i.e., there is a difference of 100 between 
the melting-point of ice and the boiling-point of water), but 
its zero is 273 below the Centigrade zero. To convert 

Nat. Portrait Cattery. 




temperatures C. into temperatures Absolute all that we have to 
do is to add 273. Thus 15 273 + 15 Absolute, or 288 
Abs. From these considerations it can easily be shown that the 
volume of a given mass of gas is directly proportional to the 
Absolute Temperature, provided that the pressure is constant. 
Suppose a gas to occupy x c.c. at C., and x' c.c. at t C., 
then from the above formula, 


x 1 273 

* ' x' t 273 + f 

1 + 273 

Suppose now that the gas occupies x" c.c. at t' C. Then 
x 273 


' 273 + t' 

273 + i 

x" 273 + t' 

or, in other words, the volume of a gas is directly proportional 
to the temperature in C. + 273, that is, to the temperature 

When correcting the volumes of gases for changes in 
temperature the first necessity, then, is to express the tem- 
perature in Absolute, by adding 273 to the temperature 
in C. A famous chemist used to put a private mark in his 
books on a certain page, for identification purposes. When 
asked upon what page he put the mark, he replied, " On 
what other page could a chemist put his mark than 273 ? " 

Charles' Law may be written shortly as 

V oc T (Absolute) if F is constant. 


Example. A gas occupies 72 c.o. at 15 C. What will be its volume 
at - 10 C. T 

Change the temperatures to degrees absolute first. 15 C. = (273 
+ 15) Abs. = 288 Abs. and - 1C* C. = (273 - 10) Abs. = 263 Aba. 
The original temperature is higher than the final, therefore the gas 


in to be cooled and therefore will contract. Hence the larger of the 
two temperature numbers must go at the bottom in the expression 

_ T . 72 x 263 

New volume = - ^-- 


= 65-5 c.c. 

72 288 

Or, by means of the formula, = ~~- 

x JQO 

:. x = 65-5 c.c. 

In all calculations it is much better to work from first 
principles than to use the formula. Never use a formula until 
you know its meaning thoroughly and can work it out for 
yourself. 1 

Normal or Standard Temperature is C. or 273 Abs. 
N.T.P. or S.T.P. are contractions for Normal or Standard 
Temperature and Pressure, C. 760 mm. 

Correction of Volumes of Gases for Temperature 
and Pressure Simultaneously. This presents no difficulty. 
Correct for Temperature first and then for Pressure, e.g., 

A gas occupies 180 c.c. at 87 C. and 1,000 mm. pressure. What 
volume will it occupy at N.T.P. ? 

Temperatures in Absolute : 87 C. = 360 Absolute. 

C. - 273 
The gas is therefore to be cooled, .*. the new volume 

180 x 273 
X ~ 360~~- 

This volume has now to be corrected for pressure. Original pressure 
1,000 mm., final 760 mm., .'. the pressure is to be lowered, .*. the 
gas will expand, .'. the new volume 

18 x 273 1*000 

~ 3(30 760 

= 179-6 c.c. 
The final volume at N.T.P. is therefore 179-6 c.c. 

A Third Gas Law may be deduced from those of Boyle 
and Charles, and is therefore nameless. It is that the pressure 

1 If you use a formula in a mechanical way you will never make 
any advance, and, to look at the matter from a very low point of 
view, any examiner can very easily set you a question in which slavish 
use of the formula will lead you hopelessly astray. To find out whether 
the candidate understands his work properly or has merely absorbed 
formulae is, indeed, one of the few cases in which an examiner may 
legitimately set traps. 


of a given mass of gas varies directly as the Absolute Temperature 
if the volume is constant, or P oc T if V is .constant. 

For, from Boyle's Law, 
and from Charles' Law 


c p 

V oc T abs. 

Combining these two expressions, V oc p ' 

T abs. 

:. P oc -y- 

and if V is constant, P oc T abs. 
Or, alternatively, the pressure exerted by a gas increases by 

- of its value at C. for every rise in temperature of 1 C., 

if the volume is kept constant. 

These three laws may be combined into one expression, 


a constant. This constant is usually called R, so that 

the gas equation becomes 

PV = RT. 

A value can be calculated for R by taking definite values for 
P, V and T. 

Suppose we take P = 76 cm. of mercury, V = 22,400 c.c* (the 
reason for this will be seen later, p. 65), and T = 273 Absolute. 
Expressing the pressure in dynes per square cm. we should have 
P = 76 x 13-5 X 981, since 13-5 is the specific gravity of mercury. 
The equation then is 

7G X 13-5 X 981 X 22,400 _ 



/. R = 8-3 X 10 7 ergs per 1 C. 

If we divide this by the mechanical equivalent of heat, 
1 calorie = 4-] 8 X 10 7 ergs, 

O.O y JQ 

we shall get the value of R in thermal measure = ^ 

= 2 calorieo. 


Correction of the Volume of a Moist Gas. If a gas IB 

collected over water, it takes up some water vapour with it. 
The gas collected therefore is not pure, but mixed with water 
vapour. In a mixture of gases it was found by D ALTON in 
1801 that each gas exerts the same pressure that it would exert 
if it were alone in the same volume. This is called the Law of 
Partial Pressures, and may be alternatively stated as " In 
a mixture of gases, the total pressure is equal to the sum of the 
partial pressures of the gases present" 

The total pressure exerted by a gas that is saturated with 
water-vapour will therefore be made up of two pressures, 
(a) the pressure of the gas itself, and (b) the pressure of the 
water-vapour. Now the pressure of water- vapour in contact 
with liquid water is found to be constant at a constant tem- 
perature. Tables have been drawn up giving the " pressure 
of aqueous vapour," or " vapour pressures of water " at 
different temperatures, and by reference to them we can at 
once find the partial pressure of the water-vapour in a ga 
saturated with the latter at a known temperature. If the 
gas is collected under atmospheric pressure, the real pressure 
of the gas itself will be the atmospheric pressure minus the 
pressure of water-vapour at the temperature concerned. 

Example. Suppose we have 100 c.o. of a moist gas at 15 C. when 
the barometer reads 770 mm. What will be its volume dry at N.T.P. f 
First, look up the pressure of water- vapour at 15 C. This is 
found to be 12-7 mm. Therefore the true pressure of the gas is 
770 12-7 757-3 mm. The rest of the correction is then made 
in the usual way. 


1. Find the volume at 300 C. of a gas that occupies 1,000 C.G* 
at 10 C. (Pressure constant.) 

2. Find the volume at 3J atmospheres of a gas that occupies 10 
litres at 740 mm. (Temperature constant.) 

3. At what temperature will a gas, that occupies 100 c.c. at 15 C., 
occupy 321 c.c., if the pressure remains constant ? 

4. Find the volume at N.T.P. of a gas that occupies 3,000 c.c. 
at 12 mm. pressure and 79 absolute. 

5. What will be the volume, dry, at N.T.P. of a gas occupying 
23'4 c.c. at 18 C. 755 mm., collected over water ? 

6. State (a) Boyle's Law, (b) Charles' Law. What is the gas equation t 


Having assumed the existence of atoms of elements and 
of " compound atoms " of compounds, Dalton's scientific 
curiosity led him on to try to determine the weights of these 
minute particles. It was clearly ridiculous to try to express 
their weights in grams, since the gram is so enormously larger 
than an atom. In the Middle Ages, the Schoolmen used 
gravely to discuss the problem of the maximum number of 
angels who could dance on the point of a needle, and con- 
eluded that there was probably room for 10,000. Atoms are 
so small that 10,000 of them would find as much room on the 
point of a needle as 10,000 people would on the continent of 
Europe ; we know now, in fact, that the weight of an atom 
of hydrogen is 0-00000000000000000000000165 grams [i.e. 
1-65 X 10 ~ 24 grams]. 

Dalton therefore contented himself with the lesser problem 
of finding the relative weights of atoms, choosing the lightest 
atom, that of hydrogen, as unit. But even this was a task 
stupendous in its magnitude, although it has now become so 
commonplace a thing that we are often apt to forget the 
sublime genius of the chemists who first successfully solved 
the problem. 

Before we proceed, it is necessary for us to get more precise 
ideas of " atoms " and " compound atoms/' so that we may 
know exactly what we propose to do. The gradual crystalliz- 
ation in the minds of chemists of the conceptions of these 
tiny particles is to be found elsewhere ; it will be sufficient 
here to give the modern definitions. 



For our purpose, then, we may take the atom to be the 
smallest, chemically indivisible, particle of an element that can 
take part in a chemical change. Physicists have shown that 
atoms are divisible, but no one has yet succeeded in splitting 
up an atom by chemical means, so that our definition is sound. 

The smallest particle of an element or compound that can 
normally lead a separate existence is called a molecule. It is 
evident that our term molecule includes Dalton's " compound 
atoms," but it also includes the ultimate particles of elements 
that can exist in a state of freedom. For example, the 
smallest particles of hydrogen that can lead a separate 
existence each consist of 2 atoms of hydrogen ; these di- 
atomic particles are called molecules of hydrogen. But the 
atoms of argon, one of the rare gases of the atmosphere, 
normally, and indeed always, lead a solitary existence. Here 
the atom of argon is also the molecule, and the molecule is 
said to be monatomic in this case. Similarly, the molecule 
of ozone is triatomic, that of phosphorus vapour is tetr atomic, 
while in sulphur vapour hexatomic and octatomic molecules 
are found. The vapours of most non-metallic elements are 
diatomic, while the vapours of metals are generally monatomic 
so far as they have been investigated. 

The smallest number of atoms that the molecule of a com- 
pound can contain is obviously two, since a compound must 
contain at least two elements, and the smallest number of 
atoms of each element that its molecule can contain is 1. 
There is no upper limit to the number of atoms the molecule 
of a substance may contain. Thus the starch molecule 
contains at least 10,000 atoms, and that of albumen probably 
a much larger number. Such molecules are said to be 

The weights, then, of atoms and molecules are expressed 
in terms of the weight of the hydrogen atom, which is taken 
as 1. Hence the Atomic Weight of an element is the number 
of times its atom is heavier than the atom of hydrogen, and th* 
Molecular Weight of an element or compound is the number 
of times its molecule is heavier than the ATOM of hydrogen. 

As in the case of Equivalents* Atomic and Molecular Weights 


are ratios, and tell us nothing of the actual weight in grams 
of the atom or molecule. All the information they give us 
is that if the hydrogen atom weighs x grams, then the atom 
of oxygen, for example, weighs 16# grams, and the molecule 
of marble 100# grams, and so on. This conception is very 
useful in practice, for if we know the relative numbers of atoms 
or molecules taking part in a reaction, and also know their 
atomic or molecular weights, we shall know also the relative 
weights in grams (or in any other units) of the substances 
concerned. Thus, suppose we knew that one molecule of 
caustic potash is neutralized by one molecule of hydrochloric 
acid, and that the molecular weight of caustic potash is 56 
while that of hydrochloric acid is 36-5, we could then say 
that when caustic potash and hydrochloric acid react they 
do so in the proportions by weight of 56 to 36-5 in any units 
we cared to choose, probably grams, but equally truly in 
pounds or ounces. 

The first step on the road to molecular weight determination 
was taken by the brilliant French chemist JOSEPH- Louis 
OAY-LUSSAC, who in 1808 formulated the Law of Gaseous 
Volumes, or as it is more generally now known, Gay- 
Lussac's Law : When gases react together their volumes are 
simply related to one another and to the volumes of the products 
if these are gaseous. 

For example, assuming all measurements to be made at the 
eame temperature and pressure, 

(i) 1 vol. of hydrogen combines with 1 vol. of chlorine to 
form 2 vols. of hydrochloric acid gas. 

(ii) 2 vols. of carbon monoxide combine with 1 vol. of 
oxygen to form 2 vols. of carbon dioxide. 

(iii) 2 vols. of hydrogen combine with 1 vol. of oxygen to 
form 2 vols. of steam. 

(iv) 1 vol. of nitrogen combines with 3 vols. of hydrogen 
to form 2 vols. of ammonia. 

Although Gay-Lussac himself did not conclude from these, 
and other results of a similar nature, that the ratio of the 
volumes of reacting gases must bear a simple relation to the 
ratio of the numbers of " atoms " in those volumes of the 


gases, the Swedish chemist BERZELXTJS saw the implication 
and somewhat rashly suggested that " equal volumes of 
gaseous elements contain equal numbers of atoms," and this- 
was afterwards extended to include all gases, not merely 
the elementary ones. Dalton and Gay-Lussac, however, both 
pointed out that this statement led to grave inconsistencies. 
Thus, 1 volume of hydrogen, containing say n atoms, will 
combine with 1 volume of chlorine, also containing n atoms 
(according to the above hypothesis), to form 2 volumes of 
hydrochloric acid. Now, if in this process 1 atom of hydrogen 
combines with 1 atom of chlorine to form 1 particle of hydro- 
chloric acid, it follows that n particles of hydrochloric acid 
will be formed. But these occupy 2 volumes, therefore 


1 volume will contain - particles, or only half as many as the 


same volume of hydrogen or chlorine. Berzelius' suggestion 
had therefore to be abandoned, but two years previously, namely 
in 1811, the Italian scientist AMEDEO AVOGADRO had already 
put forward a hypothesis which fully explained the difficulty. 

Avogadro's Hypothesis was that Equal volumes of all 
gases, under the same conditions of temperature and pressure, 
contain the same number of molecules (not atoms as Berzelius 
had supposed). 

If the Atomic Theory is the main arch of chemistry, 
Avogadro's hypothesis is the keystone of the arch. Not only 
did it clear up the difficulties which had puzzled (and even, 
it is said, created bad feeling between) Dalton, Berzelius and 
Gay-Lussac, but it gave also a ready access to the determina- 
tion of molecular weights. Yet, strange as it may seem to 
us, the hypothesis was scarcely regarded at the time and lay 
almost forgotten for nearly half a century, until this Sleeping 
Beauty was rediscovered by a Prince Charming in the person 
oi Avogadro's countryman CANNIZZARO, who in 1858 (two 
years after Avogadro's death) recalled the attention of chemists 
to the invaluable jewel they had neglected. How much 
the progress of chemistry was delayed by lack of apprecia- 
tion of Avogadro's Hypothesis it would be difficult to say. 
One has only to read some of the text- books of chemistry 


of the first half of the nineteenth century to realize the 
confusion that reigned in the minds of chemists regarding 
the true values of atomic and molecular weights. 

Let us first see how Avogadro's Hypothesis explains the 
volume changes in the reaction of hydrogen and chlorine. 
One vol. of hydrogen combines with 1 vol. of chlorine to form 
2 vols of hydrochloric acid ; that is experimental fact. 
According to Avogadro, we can say " therefore, n molecules 
of hydrogen combine with n molecules of chlorine to form 
2?i molecules of hydrochloric acid, .*. 1 molecule of hydrogen 
combines with 1 molecule of chlorine to form 2 molecules of 
hydrochloric acid." In other words, since, by the Atomic 
Theory, all the molecules of hydrochloric acid must be 
identical, the hydrogen and chlorine molecules must each have 
been halved. Now atoms are, ex hypothesi, indivisible, there- 
fore the molecules of these gases must contain an even number 
of atoms, at least two. It does not follow that they do not 
contain more than two. 

The density of a gaseous substance, that is, its mass per 
unit volume, is always very small, so chemists prefer to 
express the density of such substances in terms of the density 
of hydrogen taken as unity. This applies not only to sub- 
stances which are normally gaseous, but to the vapours of all 
substances which can be vaporized. Hence the VAPOUB 
DENSITY of a substance is the relative density of its vapour 
in terms of that of hydrogen as unity, or the number of times a 
certain volume of its vapour is heavier than the same volume of 
hydrogen under the same conditions of temperature and pressure. 

We shall arrive at an important conclusion if we consider 
this definition in the light of Avogadro's Hypothesis. We 
have, at constant temperature and pressure, 

weight of a certain volume of vapour 
* " ^ weight of the same volume of hydrogen 

weight of n molecules of the vapour 
.'. by Avogadro, V.D. = we i g ht of n molecules of hydrogen 

weight of 1 molecule of the substance 
^ weight of 1 molecule of hydrogen 

By kind pcrmitnon of R. B. Pilcher, Esq., of the Institute 
of Chemistry of Great Britain and, Ireland. 




In other words, the number that expresses the vapour density 
of a substance is also the number of times the molecule of the 
substance is heavier than the molecule of hydrogen. 

If we combine this with the definition of molecular weight, 
we shall make another important step, for 

_ _ TTT weight of 1 molecule of the substance , 

M.W. = 1.4. * i ; FTTH > and 

weight of 1 atom of hydrogen 

weight of 1 molecule of the substance 
' weight of 1 molecule of hydrogen 

Therefore the molecular weight of a substance is as 
many times its vapour density as there are atoms in 
the molecule of hydrogen. 

Now we are beginning to see our way to the determination 
of molecular weights, and can realize the importance of 
Avogadro's Hypothesis. If only we can find how many 
atoms (" x ") there are in the molecule of hydrogen we shall 
be able to find the molecular weight of any gaseous or volatile 
substance simply by measuring how many times a certain 
volume of the gas or vapour is heavier than the same volume 
of hydrogen under the same conditions of temperature and 
pressure, and then multiplying this number by " x" 

Atomicity of the Hydrogen Molecule. We have already 
found that the molecule of hydrogen must contain at least 
two atoms, and if more than two, still an even number. We 
can get further evidence from a study of the behaviour of 
acids. All acids consist of two parts, of which one is hydro- 
gen. In a great many acids all the hydrogen can be replaced 
by a metal, the products being called salts. Suppose that 
the molecule of one of these acids contains one hydrogen atom, 
and that we call the rest of the molecule X. We could then 
"represent the acid by Fig. 2. 

FIG. 2. FIG. 3. 

If we now act upon this acid in such a way as to form its 


sodium salt, it is clear that the hydrogen can be replaced in 
one stage only, for it is impossible to split an atom. The 
sodium " X-ate " would be represented by Fig. 3. 

But suppose the molecule of an acid contains two atoms 
of hydrogen ; it will then be possible to replace the hydrogen 
by sodium in two stages, forming first the compound repre- 
sented by Fig. 4, then that of Fig. 5. 


Fm. 4. Fio. 5. 

The first of these two compounds is half an acid and half 
a salt ; it is called an acid salt. The second is the normal 
sodium salt of that acid. Continuing the process, if the 
molecule of an acid contains n hydrogen atoms (let us repre- 
sent it by H W X), that acid will form n sodium salts, viz., 

Sodium \ 

Sodiurru ^ Sodium^ Sodium X\ ,, 

H n _ 1 - >X Sodmm~)X, and Sodium // 

and so on, up to Sodium") s ^ 

X or Sodium n X. 
Sodium J ' 

H w _ n 

In other words, the number of sodium salts that an acid 
will form is the number of hydrogen atoms contained in its 
molecule. Now, hydrochloric acid will form one sodium salt, 
and one only. Hence its molecule contains one hydrogen 
atom ; but we have already shown (p. 36) that its molecule 
contains half a molecule of hydrogen. 

.*. Half a molecule of hydrogen = 1 atom of hydrogen 
.". the molecule of hydrogen contains 2 atoms. 
Probably while you have been reading it you have had an 
uneasy feeling that all is not well with this argument, and 


you may even have arrived at some definite criticisms of it. 
See if they agree with the following 

1. How do we know that hydrochloric acid is one of those 
acids all of whose hydrogen is replaceable by a metal ? The 
answer to this criticism is that no one has yet succeeded in 
preparing any sodium chloride containing hydrogen. (This 
can easily be tested by analysis.) 

2. Are we not pledging the future by assuming that, because 
no one has yet prepared an acid sodium chloride, no one ever 
will ? Here the answer is that until some one does so, we are 
entitled to believe that it is impossible, especially after a 
hundred and fifty years of systematic chemistry. 

3. How do we know that one atom of sodium will replace 
one atom of hydrogen ? This seems to be a much more 
formidable problem, for there seems no valid reason a priori 
why one atom of sodium should not be capable of replacing 
say two atoms of hydrogen. Suppose, for the sake of argument, 
that it does. How is the " proof " of the diatomicity of the 
hydrogen molecule affected ? If the acid contains one atom 
of hydrogen per molecule, it is not affected at all, for the 


sodium salt must be [ sodium | <V ; no other is possible. If 

the acid molecule contains two atoms of hydrogen, the follow- 
ing sodium salts would be possible 

H \ 


Sodium X, Sodium^ 

> X 

but so would salts of the type 


/ X 


Sodium/ etc. 



However, these facts are really irrelevant, since however 
many atoms of hydrogen one atom of sodium will replace, it is 
obvious from the above that an acid containing one atom of 
hydrogen per molecule will form only one sodium salt, while 
any acid whose molecule contained more than one atom of 
hydrogen could always form more than one sodium salt. We 
are therefore justified in assuming that if hydrochloric acid 
will form only one sodium salt, its molecule contains only one 
atom of hydrogen. 

It must, however, be admitted that this proof of the 
diatomicity of the hydrogen molecule is not altogether satis- 
fying, especially for a fact of such basic importance. It rests 
on negative, results, namely, the impossibility of preparing 
acid chlorides, and can therefore never be as satisfactory as a 
positive proof. Fortunately, other evidence is forthcoming, 
and although a strictly logical proof cannot be given in this 
place, the following confirmatory facts may set at rest the 
mind of a doubting Thomas. 

(i) In the reactions of hydrogen with other gases, it is 
common to find that one volume of hydrogen gives two volumes 
of the gaseous product, but never more. If the molecule of 
aydrogen contained more than two atoms, we might expect that 
occasionally one volume of hydrogen would give more than 
two volumes of a gaseous product. This is negative evidence 
again, but is nevertheless extremely significant. 

(ii) Considerations of the specific heats of gases lead to the 
conclusion that the molecule of hydrogen is diatomic. The 
specific heat of a substance is the number of calories required 
to raise the temperature of 1 gram of a substance through 
1 C. (A calorie is the amount of heat required to raise the 
temperature of 1 gram of water through 1 C. Be careful to 
avoid saying " to raise 1 gram of the substance through 
1C." it is not the substance that you "raise," but its 
temperature.) Now if a gas is heated, it expands, and in doing 
BO will perform work, as in the expansion of the gases in the 
cylinder of a motor. The energy necessary for the perform- 
ance of this work is taken from the heat supplied, and to raise 
the temperature of 1 gram of a gas through 1 C. will therefore 



require more heat if the gas be allowed to expand and do 
external work than if the gas be kept at constant volume and 
hence prevented from doing external work. A gas therefore 
has two specific heats, the specific heat at constant volume, and 
the specific heat at constant pressure (when it is allowed to 
expand) ; and the latter will be the greater. Now if the 
molecules of a gas are monatomic, all the heat supplied to the 
gas will go to increase the speed of translation of the molecules, 
that is, the speed with which they move about ; but if the 
molecules of the gas contain more than one atom, part of the 
heat supplied will go as before to increase the speed of trans- 
lation of the molecules, but a part will be used in increasing 
the speed of vibration of the atoms within the molecule. It can 
be shown theoretically that in the latter case the ratio 

specific heat at constant pressure (C^) , . , 

- . x *f. 18 always less than for 

specific heat at constant volume (CJ J 

a gas whose molecule is monatomic, and the greater the 


number of atoms in the molecule the smaller the value of ~ . 

C t 


For a monatomic gas, -^ = 1-67 ; diatomic gases give the 
O r 

approximate value of 14, and so on. The value for hydrogen 
is 1408. 

(iii) No hydrogen compound yet discovered contains less 
hydrogen than that found in an equal volume (as gas) of 
hydrochloric acid. 

(iv) No facts have yet been discovered that do not har- 
monize with the assumption that the molecule of hydrogen 
contains two atoms. 

Note, however, that not one of the above arguments is a 
logical proof that the molecule of hydrogen is diatomic. The 
diatomicity of hydrogen has been conclusively proved by 
physical means, but even if it had not, the above evidence 
would justify us in assuming it until something turned up to 
disprove it. 

We may therefore now say that the Molecular Weight 
of a substance is twice its Vapour Density, since the 
molecule of hydrogen contains two atoms (see p. 38). 


Methods of Determining Vapour Densities. 

A. Gases. The principle of this method is very simple, 
but to get accurate results many precautions have to be taken 
and corrections made. A glass or metal globe, fitted with a 
stopcock, is evacuated by means of an air-pump and then 
weighed (w gms.). It is then filled with the gas and weighed 
again (m gms.). Weight of gas = m w gms. The globe 
is now evacuated once more and then filled with hydrogen 
and weighed (n gms.). Weight of hydrogen = n w gms. 

/. V.D. = if the temperature and pressure have 

remained constant throughout. 

FIG. 6. Regnault's Vapour Density Apparatus. 

In accurate work, great care has to be taken to see that aD 
the measurements are carried out at exactly the same 
temperature and pressure, or at known temperatures and 
pressures if they are different, when a correction has to be 
made according to the laws of Boyle and Charles. It is also 
important to note that the real weight of the globe, that is, 
its weight in vacuo, is its apparent weight in air + the weight 
of air displaced. Although the difference between the real 



weight and the apparent weight is very small, it is sufficient 
to affect the results if not allowed for. Moreover, the weight 
of air that the globe displaces will depend upon the tem- 
perature, pressure and relative humidity of the air, and 
owing to the water- vapour present in the air there is always 
a film of moisture on the surface of glass, porcelain and metal. 
To make the necessary corrections for all these things would 
be very tedious and, probably, the corrections would not be 
very accurate when made. Regnault solved the problem by 
using a second similar globe which he hung on the other arm 
of the balance. The errors due to the above conditions now 
balanced one another, so that the corrections were unnecessary. 
A final correction, suggested by Lord Rayleigh, deals with the 
shrinkage in volume of the globe on 
evacuation. This has to be found by 
experiment, and allowed for. 

B. Volatile Liquids (and Solids). 
(i) Victor Meyer's Method. Victor 
Meyer and his pupils used to make 
large numbers of new compounds, and 
wanted a method of determining 
their vapour densities (and thence 
their molecular weights) which should 
be simple, quick and reasonably 
accurate. Meyer therefore devised the 
following method, which is much more 
convenient than that of Dumas 

The apparatus consists of an inner 
tube A, the lower end of which is 
enlarged into a cylindrical bulb, while 
the upper end is enlarged to carry a 
rubber stopper H. Towards the upper 
end of the tube A there is a side-neck 
B communicating with a pneumatic 
paratus. trough C . The inner tu be is surrounded 

by a wider outer tube or jacket G, at the bottom of which 
is a bulb F in which are placed 70-80 c.c. of a liquid whose 

FIG. 7. Victor Meyer 
Vapour Density Ap- 


boiling-point is at least 30 higher than that of the liquid 
whose vapour density is required. The apparatus is set up 
as shown, except that the graduated tube I) is not yet placed 
over the end of the side-neck. The liquid in F is boiled and 
the heating continued until the temperature of the inner tube 
is constant, that is, when no more bubbles of air escape from 
the end of the side-neck and no water is sucked back. When 
this is so, the graduated tube, filled with water, is placed in 
position (as is shown in the figure). 

Meanwhile a small bottle, of such a size that it slips easily 
down the tube A, is weighed empty and is then filled with 
the liquid whose V.D. is required, and weighed again. Next, 
the cork is removed, the weighed bottle and contents dropped 
in, and the cork replaced all as quickly as possible. The 
bottle falls to the bottom of the inner tube, where a little 
mercury may be placed to break its fall. The liquid in the 
small bottle now finds itself in a place which is at a tempera- 
ture some 30 higher than its boiling-point. It therefore 
vaporizes very rapidly and blows the stopper out of the bottle ; 
and the vapour, advancing up the tube en masse, drives its 
own volume of air over into the graduated tube D. The 
graduated tube is then transferred to a deep jar full of water, 
in which it is lowered until the levels of the water inside and 
outside the tube are the same, and the volume of the air read. 
This process of " levelling " is necessary in order to get the 
moist air in the tube at atmospheric pressure. 

The height of the barometer is then read, and the tempera- 
ture of the water in the pneumatic trough taken. This 
temperature will be the temperature of the moist air in the 
graduated tube, since the air has passed through the water. 
(Of course, if the volume of air is not read off at once, it would 
be more accurate to take the temperature of the air surround- 
ing the tube.) We now have the following data 

Weight of bottle + liquid = x gms. 
Weight of bottle = y gms. 
/. Weight of liquid = x y gms. 

Barometer = say 770 mm. Temperature = say 15 C* 


Volume of moist air after levelling = m c.o. 

1 litre of hydrogen at N.T.P. weighs 0-09 gms. 

From these data we can calculate the -Vapour Density of 
the liquid. All gases obey the same gas-laws, therefore the 
volume of air in the graduated tube is equal to the volume which 
x y gms. of the substance in the state of moist vapour would 
occupy if it could exist under the same conditions of temperature 
and pressure. The whole point of the method is the ingenious 
way in which the vapour of the liquid is made to drive over 
its own volume of air, which remains gaseous under conditions 
in which the vapour itself could not. The reason for having 
the temperature of the outer jacket so much higher than the 
boiling-point of the liquid under experiment will now be 
clear : it is to ensure that the vapour of the liquid shall come 
off with a rush, and therefore drive the air over before diffu- 
sion of the vapour into the air can take place. If diffusion 
occurred, the gas driven through the side-neck would consist 
partly of air and partly of vapour, and as the latter would 
condense in the water of the trough, the experiment would 
be useless, since the volume of air collected would no longer 
be equal to the volume of vapour produced. The shape of 
the inner tube also helps to prevent diffusion. 

The result is calculated as follows : 

Volume of moist air at 770 mm. 15 C. = m c.c. 

Vapour pressure of water at 15 C. = 12-7 mm. 

.". true pressure of the air = 770 12-7 mm. = 757-3 mm. 

Temperature = 15 C. = 288 Absolute. 

TT i * j xivTrn^ mX 757*3x273 

/. Volume of air, dry, at N.T.P. = - 76Q x 2gg 

= n c.c. 

.'. n c.c. of the vapour at N.T.P. would weigh x y gms. 
But n c.c. of hydrogen at N.T.P. weigh n x 0-00009 gms. 

and the Molecular Weight of the substance will be twice 

For liquids that boil up to 70 C., water may be used in 
the outer jacket. For those with higher boiling-points suit- 


able substances to use in the outer jacket are aniline, 
B.P. 182 ; nitrobenzene, B.P. 208 ; sulphur, B.P. 444, etc. 
(ii) Dumas 9 Method. A glass globe provided with a long 
neck drawn out to a point (see Fig. 8) is weighed, and then 
a few grams of the liquid whose V.D. is required are intro- 
duced. This is done by dipping the end of the neck into some 
of the liquid in a crucible and then gently warming the globe. 
The air inside expands, and some is driven out, so that on 
allowing the globe to cool a small quantity of the liquid is 
drawn in. The globe is then immersed in a bath the tern- 

FIG. 8. Dumas Vapour Density Globe. 

perature of which is constant and at least 30 above the 
boiling-point of the liquid. This temperature is recorded. 
When vapour ceases to be expelled from the globe (which is 
easily ascertained by holding a clean piece of glass opposite 
to the neck when the glass is no longer dimmed, the vapour 
has stopped coming out), and the last drop of liquid has vola- 
tilized, the end of the neck is sealed in the blowpipe flame. 

The globe, which is now full of vapour, is removed fron? 
the bath, dried, cooled and weighed. The end is then broken 
off under freshly-distilled (i.e., air-free) water, when the globe 
fills with water. It is dried on the outside and weighed 


again, together with the end of glass broken ofi. Generally 
a small bubble of air is left in the globe ; this is air which the 
vapour has not expelled, and a correctfon must be made 
for it. 

The capacity of the globe is found by subtracting the weight 
of the globe full of air from its weight when filled with water, 
since the weight of the air may be neglected in comparison 
with the weight of the water. The weight of the water 
in grams is therefore equal to the capacity of the globe 
in c.c. 

From this volume, the weight of the air filling the globe 
at the observed temperature and pressure can be calculated, 
since 1 c.c. of air at N.T.P. weighs 0-00129 grams. The 
weight of the globe filled with air minus the weight of the air 
gives the weight of the globe empty. Subtracting this from 
the weight of the globe full of vapour gives the weight of the 
vapour. Thus we know the weight of vapour which occupies 
a known volume at known temperature and pressure, and 
hence the vapour density can be calculated, since 1 litre of 
hydrogen at N.T.P. weighs 0-09 grams. 

Although this method is capable of giving more accurate 
results than that of Victor Meyer, it is much more tedious 
to carry out and more troublesome in calculation, and hence 
is rarely used except in special cases. 

The result may be worked out as follows : 

Weight of globe filled with air at temperature t and pressure p = x gms. 
,, ,, ,, ,, vapour ,, t' and pressure p = y gms. 

, f ,, ,, ,, water z gms. 

.*. Capacity of globe in c.c. = weight of water in grams 

= z x. 
z x c.c. of air at t C. p mm. pressure become 

/. weight of the z x c.c. of air at t C. p mm. pressure 

/_ _ \ O'TO fft 

" (273 4- *) 760 X ' 00129 S 1 ^ 1118 ' since l c - of air at N.T.P. weighs 
00129 gms. 

.-. weight of globe empty - x - [{* ~ f** X 0-00129] gma. 



m grams 

.'. weight of x c.c. of vapour at t' and p mm. = y m grams. 
If y fn grams of vapour occupy z x c.c. at t' and p mm., they 
will occupy 

/2 ,\ v 273 X f> 

(273 -f- *') X 760 C ' C * at N ' T - P - 

= n c.c. 
But n c.c. of hydrogen at N.T.P. weigh 0-00009 X n gms. 

.'. Vapour Density = ^ ~ m 

^ J n X 0-00009 

and the molecular weight will be twice this. 

(iii) Hofmann's Method. Certain sub- 
stances decompose when heated to their 
boiling-points under atmospheric pres- 
sure, but will vaporize without decom- 
position under lower pressures. It is 
evident that Victor Meyer's method and 
Dumas' method are inapplicable here, and 
Hofmann's method is used in these cases. 
The principle is to vaporize the substance 
at a temperature and pressure such 
that it is not decomposed, and then to 
find the volume of a given weight of 

The apparatus consists of a wide baro- 
meter-tube which is graduated on one 
side from the bottom upwards in milli- 
metres, and on the other from the top 
downwards in cubic centimetres. This 
tube is filled with mercury and inverted 
in a trough of mercury, thus giving rise 
to a Torricellian vacuum. 

The tube is now surrounded by a jacket 
through which the vapour of a liquid 
that boils at a suitable temperature Fia. 9. Hofmann 
may be passed. The vapour passes out 
at the bottom of the jacket and may 

Vapour Density 

be condensed, if desired, by means of a Liebig condenser. 
A small quantity of the liquid whose vapour density is 


required is weighed out in a tiny bottle similar to that 
used in the Victor Meyer method, and this bottle is introduced 
into the barometer tube, when it rises to the top of the mer- 
cury and enters the vacuum. Here, owing to the low pressure, 
the liquid volatilizes at the temperature of the vapour in 
the jacket, which may be much below that of its boiling- 
point under atmospheric pressure. 

The vapour of the boiling liquid in the flask is passed through 
the jacket for some time until the level of the mercury, 
which has fallen, remains constant. The height of the mer- 
cury column is then read off on the millimetre scale, and the 
difference between this height and the original height of the 
mercury gives the pressure on the vapour. The temperature 
is taken as that of the boiling-point of the liquid in the 
flask. The volume of the vapour is read from the c.c. 

Weight of liquid (and therefore of vapour) = m gms. 
Original height of mercury column = a mm. 

Final ,, ,, ,, ,, =6 mm. 

.'. pressure of vapour (a b) mm. 

Temperature = t C. Volume of vapour = v c.c. 
Now v cc. at t C. and (a 6) mm. pressure become 

*f 3 (*-*) c . c . at N.T.P. 

(273 + 0-760 
= d c.c. 

But d c.c. of hydrogen at N.T.P. weigh d X 0-00009 gms. 

d X 

and the molecular weight will be twice this. 

For other methods of determining Molecular Weights see Chapter 


1. Define atom, molecule, atomic weight, molecular weight. 

2. State Avogadro's Hypothesis and explain its importance in 

3. Suppose that Avogadro's Hypothesis were proved to be untrue. 
What do you think would be the effect upon chemistry ? 

4. Define vapour density. What is the relationship between the 
molecular weight and the vapour density of a substance ? 



5. Prove, as far as you can, that the molecule of hydrogen contains- 
two atoms. 

6. Describe Victor Meyer's method of determining vapour densities. 

7. Describe Dumas' and Hofmann's methods of finding vapour 
densities. Under what circumstances is Hofmann's method especially 
useful ? 


In the last chapter we saw how it was possible to determine 
the molecular weights of those elements and compounds 
that can be obtained in the state of gas or vapour. Our 
next task is to find atomic weights, that is, those numbers 
which express the ratio of the weight of an atom of an 
element to the weight of an atom of hydrogen. It so happens 
that the molecules of many elements that are normally 
gaseous consist of two atoms, hence in these cases, and purely 
by chance, the vapour density, being half the molecular weight, 
is numerically equal to the atomic weight of the element. 
This occurrence misled the chemists of the first half of the 
nineteenth century into the error of supposing that the vapour 
density of an element was always equal to its atomic weight. 
It is clear, however, that this equality occurs only in those 
oases where the molecule of the element is diatomic. If, for 
example, the molecules of an element are monatomic, as is 
the case with argon and the other gases of the argon group, 
the molecular weight and atomic weight will be identical, 
and therefore the vapour density will here be half the atomic 
weight. In short, from the vapour density alone it is impos- 
sible to determine the atomic weight of an element. Other 
methods are fortunately not lacking. 

(i) The smallest number of atoms of an element which can 
be present in the molecule of any compound of that element 
is 1, since atoms are indivisible. That is, the smallest weight 
of any element which can be present in the molecular weight 
of any of its compounds is the atomic weight of that element. 




Note that it does not follow that the smallest weight of the 
element which is present in the molecular weight of any of 
its compounds is necessarily the atomic weight of the element, 
since it is possible to conceive that certain elements might not 
form compounds containing less than 2 or 3 orn atoms of those 
elements per molecule. Here again, however, we are justified 
in taking as the atomic weight of an element the smallest 
weight of that dement which is found in the molecular weight 
of any of its compounds, until we find some fact or other 
that is not in accordance with the value we have obtained. 
If the values we get for atomic weights are supported by 
other evidence, and nothing is discovered which shows that 
these values are incorrect, then we may assume that they are 
the true values. 

Our first method, then, involves two processes (a) the 
preparation of a great many compouDds of the element con- 
cerned, and the determination of their molecular weights, 
and (6) quantitative analysis of these compounds in order to 
determine the weight of the element present in the molecular 
weight of each. The smallest weight of the element that 
is ever found in the molecular weight of any of its compounds 
is then taken to be the atomic weight of that element. 

In general, it will be most convenient to take volatile 
compounds of the element, since the molecular weights of 
these will be easily determinable by vapour-density methods. 

As an example, let us consider the case of carbon. This 
element forms many volatile and gaseous compounds, some 
of which give the following data by the above method : 

Acetylene . 
Ethylene . 
Methane . 
Carbon monoxide 

From these figures it is concluded that the atomic weight of 
carbon is 12. 

Vol. Wt. 

Parts by 
Wt. of Carbon 
in Mol. Wt. 














If, at some future date, a carbon compound were discovered 
which contained only 6 parts by weight of carbon in its 
molecular weight, we should then have to* admit that the 
number 12 as the atomic weight for carbon is incorrect. But, 
as in many other cases in chemistry, collateral evidence con- 
cerning atomic weights can be obtained, and when all the 
evidence for any one atomic weight points to the same con- 
clusion we may feel confident that this conclusion is correct. 

Dulong and Petit's Law (1819) has proved a valuable 
instrument for atomic weight determinations. These two 
chemists found that in a great many cases the atomic weight 
of a metal X its specific heat = 64 approximately. Hence, if 
we wish to find the atomic weight of a certain metal, we could 
determine its specific heat and divide it into 6-4 ; the quotient 
would give us an approximate value for the atomic weight. 
It may be asked of what use an approximate value for the 
atomic weight is ; surely we want to know atomic weights 
as accurately as possible ? To answer this question we shall 
have to consider the question of valency. 

Valency. The subject of valency is a very large one, and 
whole books have been written on it it has, in fact, been 
defined as " that subject of which freshmen at the University 
think they know everything and of which mature chemists 
realize that they know nothing." Fortunately for our imme- 
diate purpose we do not need "to study it in detail, and can 
easily arrive at a definition of valency that will meet our 

The atoms in a molecule are not merely lying side by side ; 
they are bound together by chemical force. It is the question 
of the nature of this force that introduces the difficulties 
of valency ; here, therefore, we will simply accept it as a 
fact, and will leave the closer discussion of it for a more 
advanced course. 

^ Chemists have found, however, that the power which any 
atom of a given element has of combining with atoms of 
hydrogen is always the same. Thus any atom of oxygen is 
capable of combining with 2 atoms of hydrogen ; an atom 
of chlorine can combine with 1 atom of hydrogen ; that of 


phosphorus with 3 of hydrogen ; and that of carbon with 4 
of hydrogen. No element is known which will combine with 
hydrogen in such a way that there is a smaller proportion 
of hydrogen in the compound than 1 atom of the element 
to 1 of hydrogen ; in other words, there is no known ele- 
ment of such small combining power that 2 or more atoms of 
it are required to hold 1 atom of hydrogen in combination. 

For this reason, chemists have taken as the unit of combining 
power of an atom the ability to hold in combination 1 atom 
of hydrogen. The number of atoms of hydrogen with which 
1 atom of an element will combine is called the VALENCY 
of that element. 1 Some elements, however, will not combine 
with hydrogen, but will do so with other elements. It is 
evident that in such cases the valency of the element cannot 
be directly measured, but if we remember that ability to hold 
in combination 1 atom of hydrogen is merely the " unit of 
combining power," we shall be able to measure the valency 
in an indirect way. For example, zinc will not combine with 
hydrogen, but it will do so with oxygen, in such a way that 
1 atom of zinc combines with 1 atom of oxygen. Now 
1 atom of oxygen will combine with 2 atoms of hydrogen ; 
we therefore assume that if zinc and hydrogen could be made 
to combine they would do so in the proportion of 1 atom 
of zinc to 2 of hydrogen, and therefore say that the valency; 
of zinc is two. Similarly, gold will not combine with hydrogen^ 
but will combine with chlorine in such a way that 1 atom of 
gold combines with 3 of chlorine. Since 1 atom of chlorine 
will combine with 1 of hydrogen, we say that the valency of 
gold is 3. 

Atoms are chemically indivisible, therefore the valency of 
an element must always be a whole number. 

We shall now be able to see that there is an intimate 
relationship between the three numbers that represent 
respectively the valency, atomic weight and equivalent 
of an element. Suppose that 1 atom of an element,J 
atomic weight, say, 20, combines with 1 atom of hydm^n. 

1 This, of course, is only one of the conceptions ^f^jUency. 
There are many others. 


We could represent the molecule of the compound so formed 
(Fig. 10) as 

FIG. 10. 

What is the equivalent of this element ? We know that the 
equivalent of an element is the number of units of weight of 
it that will combine or otl erwise react with 1 of the same 
units of weight of hydrogen. Let the unit of weight chosen 
in this case be the weight of 1 atom of hydrogen. Then since 
the atomic weight of the element is 20, 20 units of weight 
of it have combined with 1 unit of weight of hydrogen, 
/. its equivalent = 20. But the atomic weight of the element is 

, nr . . A1 . Atomic Weight .. T> . * . 

also 20, . . in this case ^ : = ? = 1. But 1 atom 


of the element has combined with 1 atom of hydrogen, 
therefore the valency of the element 1, 

Atomic Weight TT , 

/. here, at any rate, ^ : = ^ = Valency. 
J Equivalent J 

Now let us suppose that 1 atom of the element combines 
with 2 atoms of hydrogen ; the molecule will be as in 
Fig. 11. 

FIG. 11. 

What is the equivalent of the element here \ Choosing 
the same unit as before, namely, the weight of 1 atom 
of hydrogen, it is clear that 20 units by weight of the 
element have combined with 2 units by weight of hydrogen, 

/. the equivalent of the element = - = 10. The atomic 


.,,.,, , , ft - Atomic Weight 20 n 
weight of the element = 20, /. ^ : -, ^ = r- = 2. 
& Equivalent 10 


But 1 atom of the element has combined with 2 of hydrogen, 
therefore the valency is 2, and here again 

Atomic Weight TT , 
: - 2 = Valency. 

Similarly, if 1 atom of the element combines with n atoms 
of hydrogen, a figure will show that the equivalent 

Atomic Weight. Atomic Weight TT . rrll . 

= 1. or . . = Valency. This is a 

n Equivalent J 

very important conclusion, for the equivalent of an element 
may easily be determined with a high degree of accuracy, and 
we know from the above that if we can find the valency (which 
must be a whole number) we shall be able to get the atomic 
weight, since atomic weight = equivalent X valency. 

It is here that Dulong and Petit 's Law becomes of value. 
Suppose that you were fortunate enough to discover a new 
metal ; how would you determine its atomic weight ? There 
is still (1931) one metal or so to be discovered, so the question 
is not of mere theoretical interest ! 

The first thing to do would be to obtain the metal in a 
state of purity. When this has been done, the next step is 
to determine the equivalent as accurately as possible by one 
of the methods previously described. If we knew the valency 
we could obtain the atomic weight by multiplying the 
equivalent and valency together. The next stage, therefore, 
is to find the valency. This is done by means of Dulong and 
Petit's Law. The specific heat of the metal is found by the 
method of mixture or by means of Bunsen's Ice Calorimeter, 
and if the number so obtained is divided into 64 the quotient 
will be the approximate atomic weight. If we now divide the 
approximate atomic weight by the equivalent, we shall get 
the approximate valency. But the valency must be a whole 
number, therefore the nearest* whole number to the approxi- 
mate valency is taken to be the true valency. This multi- 
plied by the exact equivalent will give the true atomic 

What if the approximate valency came out to, say, 2-5 ? 
Which is the true valency, 2 or 3 ? The procedure in such 


circumstances is first to feel annoyed and immediately after 
to tackle the problem in another way. 

MITSCHEBLICH, in 1819, from a study of the shape of the 
crystals of a great many compounds, came to the conclusion 
that compounds of the same class that have analogous con- 
stitutions crystallize in the same form. This statement is some- 
times known as the Law of Isomorphism, or Mitscher- 
lich's Law. (Two substances are said to be isomorphous 
when they crystallize in the same form.) 

Applied in the reverse way, this law is useful in the deter- 
mination of atomic weights, for it is assumed that compounds 
of the same class that crystallize in the same form have 
analogous constitutions. For example, suppose we were 
investigating the atomic weight of chromium, and that we 
had already determined its equivalent. We should then be 
trying to find its valency. Now when potassium sulphate 
and aluminium sulphate, in the proportion of 1 molecule of 
one to 1 of the other, are dissolved in water together, and the 
solution evaporated to crystallization, they combine together 
and crystallize out as a compound called alum, which forms 
beautiful crystals whose form can be very easily measured. 
If potassium sulphate and chromium sulphate are mixed in 
solution and the solution then evaporated to crystallization, 
it is possible to get crystals of a " chrome alum " which are 
isomorphous with those of ordinary alum. According to 
Mitscherlich's Law, we should be able to conclude that since 
chromium has replaced aluminium in alum, giving rise to an 
isomorphous chrome alum, chromium and aluminium have 
the same valency, otherwise the constitutions of these two 
isomorphous compounds of the same class would not be 
analogous. Hence, if we know that the valency of aluminium 
is 3, we may justifiably conclude that the valency of chromium 
is 3, and from this and the equivalent weight we can get the 
atomic weight. 

The Periodic System also is useful in fixing atomic 
weights. (See Chapter XV.) 

In the case of those gases or vapours the ratio of whose 
specific heats shows them to consist of monatomic molecules, 


it is, of course, possible to get their atomic weights from their 
vapour densities only. For their molecular weights = their 
vapour densities X 2, and since the molecules are monatomic 
their molecular weights their atomic weights. 

To sum up, we may say that in most cases atomic weights 
are determined by (i) finding the equivalent weight, then (ii) 
obtaining as much evidence as possible to show what multiple 
of the equivalent weight the atomic weight is. 


1. Explain why the determination of molecular weight generally 
has to precede the determination of atomic weights. 

2. State Dulong and Petit's Law. Of what use is it ? 

3. Explain shortly what you understand by valency. 

4. Prove that the atomic weight of an element is equal to its equiva- 
lent multiplied by its valency. 

5. Suppose you were to discover a metal that is still unknown. 
How would you determine its atomic weight ? 

6. State the Law of Isomorphism and explain its application. 

7. 1 -081 gms. copper displace 3-670 gms. silver from a solution of silver 
nitrate. Find the equivalent of copper, given that that of silver 
= 107-88. 

8. 0-4717 gms. zinc when dissolved in sulphuric acid give 171 c.c. 
hydrogen, collected over water, at 12 C. and 760-5 mm. pressure. 
Calculate the equivalent of zinc. 

9. 2-5537 gms. copper oxide heated in a current of hydrogen gave 
2-0391 gms. copper. Find the equivalent of copper. (Oxygen equi- 
valent = 8.) 

10. 0-21 78 gms. magnesium, when dissolved in hydrochloric acid, gave 
218-2 c.c. hydrogen collected over water at 17 C. and 754-5 mm. 
pressure. The specific heat of magnesium is 0-245. Find the atomic 
weight of magnesium. 

11. 4-323 gms. nickelous chloride are found to contain 1-956 gms. 
nickel. The specific heat of nickel is 0-1092. Find the atomic weight. 
The equivalent of chlorine is 35-5. 

12. The iodide of a metal is found to contain 55-86 per cent, iodine. 
The specific heat of the metal is 0-033. Find its atomic weight. The 
equivalent of iodine is 126-92. 


It is convenient to use symbols to represent atoms and 
molecules. DALTON used a symbolic notation of circles 
plain, containing a dot, containing a cross, blacked in, and 
so on to represent the atoms of various elements ; " com- 
pound atoms " [molecules] were represented by groups of 
these circles. Some of Dalton's symbols are shown in Fig. 
12. A " compound atom " of water (which he believed to 


Oxygen. Nitrogen. Hydrogen. Carbon. 
FIG. 12. 

consist of one atom of hydrogen and one of oxygen) was 
represented as follows : 



FIG. 13. 

This system, however, had many drawbacks and in 1811 
BBRZELITJS suggested the system that we now use. He repre- 
sented an atom of an element by the initial, or initial and 
another characteristic letter, of the name of the element, 
sometimes using the Latin (or pseudo-Latin) name in 
cases where two or more elements had names with the 
same initial letter. The following table gives some 
examples : 






Phosphorus . 


Hydrogen . 

Potassium . 


Iron . 















For a complete table of symbols and atomic weights see 
p. 674. 

A molecule of hydrogen consists of 2 atoms ; the formula 
for this is ILj. 2H stands for 2 atoms of hydrogen not com- 
bined together. 2H 2 stands for 2 molecules of hydrogen each 
consisting of 2 atoms. 30 2 means 3 molecules of oxygen each 
consisting of 2 atoms, while 20 3 means 2 molecules of ozone 
each containing 3 atoms of oxygen. 

The formulae for compounds are written similarly. Thus 
H 2 S0 4 stands for 1 molecule of sulphuric acid, consisting of 
2 atoms of hydrogen combined with 1 of sulphur, and 4 of 
oxygen. If we wish to represent the way in which the atoms 
of a molecule are arranged, we use structural or graphic 
formulae. The molecule of alcohol is C 2 H 6 0, but this tells us 
nothing of the way in which the atoms are arranged ; if, 
however, it is important to express this we write the formula 
in the following way : 

H H 

H C C H, 

I I 
H H 

where the lines represent " valencies " or unit combining 
powers. Such a formula evidently gives us much more 
information than an ordinary formula. 

It is most important to note that H 2 0, for example, does 


NOT stand for " water." The symbol Id. does not stand for 
" wealth," but for a single, minute and practically indivisible 
particle of wealth. In exactly the same way, H 8 stands for 
one MOLECULE of water, and not for " water " in general. 
If, therefore, you say " I dissolved the substance in H 2 0," 
you are, actually, saying that you dissolved the substance 
in 1 molecule of water, which you probably do not mean. 
It is, of course, permissible to use formulae instead of names 
in writing rough notes, and so on, but it is incorrect and should 
always be avoided if possible. 
The Professor, indeed, who now is 

" no more, 

For what he thought was H 2 O 
Was H 2 SO 4 " 

must have died from some other cause, since it is difficult to 
imagine a constitution so delicate that 1 molecule of sulphuric 
acid could upset it ! 

In the same way that formulae are used to represent atoms 
and molecules, equations are used to represent reactions 
between atoms and molecules. What information, for 
example, does the equation 

Zn + H 2 S0 4 = ZnS0 4 + H a 

give us ? It tells us 

(i) That under certain conditions not specified, zinc and 
sulphuric acid will react together in such a way that hydrogen 
is displaced from the acid and zinc takes its place. 

(ii) That 1 atom of zinc will react with 1 molecule of 
sulphuric acid, containing 2 atoms of hydrogen, 1 of sulphur 
and 4 of oxygen, to give 1 molecule of zinc sulphate, con- 
taining 1 atom of zinc, 1 of sulphur and 4 of oxygen, and 
1 molecule of hydrogen, containing 2 atoms. 

(iii) That the valency of zinc is 2, and that the valency 
of the S0 4 group of atoms is 2. 

If we know the atomic weights of the elements concerned 
(Zn = 65 ; H = 1 ; S = 32 ; O = 16) it tells us 

(iv) That 65 units of weight of zinc will react with 98 units 
of weight (= 2 + 32 + 64) of sulphuric acid to give 161 




units of weight of zinc sulphate and 2 units of weight or 

What information about the reaction does it not give us, 
information that perhaps we should like to have ? 

(i) It does not tell us whether the reaction is accompanied 
by evolution or absorption of heat. 

(ii) It tell us nothing of the physical states of the substances 

(iii) It does not tell us whether it is necessary to dilute the 
acid or not. 

(iv) It does not tell us whether it is necessary to apply 
heat or not. 

(v) It tells us nothing of the time taken by the reaction. 

(vi) It does not tell us whether the action is reversible or 

An equation, therefore, gives us a good deal of information 
about the reaction which it represents, but leaves almost as 
much untold. An equation, however, is very useful since the 
information that it docs give is definite and is also that which 
as a rule it is most important to know. We shall see, too, from 
the following considerations, that if gases appear in a reaction, 
the equation tells us what volume of those gases will be used 
or formed for a known weight of material. 

It is found by experiment that 2 grams l of hydrogen at 
N.T.P. occupy 22-4 litres, i.e. n molecules of hydrogen at 
N.T.P. occupy 22-4 litres. Now 22-4 litres of any other 
gas at N.T.P. would also contain n molecules, by Avogadro's 
hypothesis ; therefore the ratio 

weight of 22-4 litres of a gas at N.T.P. __ 
weight of 22-4 litres of hydrogen at N.T.P. 

weight of n molecules of the gas the number representing 

weight of n molecules of hydrogen^ tne vapour density (A). 
But in this case, n is the number of molecules of hydrogen 
that weigh 2 grams, .'. the weight in grams of n molecules 
of the gas must be the molecular weight in grams, to get the 
necessary result (A) that 

1 Actually 2-016 grams, since the standard of atomic weights is now 
taken as O 16, upon which H = 1-008. The difference may be 
neglected at this stage. 


weight of n molecules ( M.Wt. in grams) of the gas 
weight of n molecules of hydrogen (= 2 grams) 

should give the number representing the vapour density. 

In other words, the molecular weight in grams (sometimes 
called the G.M.W. or Gram-Molecular- Weight or Gram-Mole- 
cule) of any gas at N.T.P. occupies 22-4 litres (sometimes 
called the G.M.V. or Gram-Molecular-Volume). By means of 
the gas laws it is of course now possible to calculate what 
volume the G.M.W. of a gas would occupy at any other 
temperature and pressure. 1 

Hence, in the equation 

CaCO 3 + 2HC1 = CaCl a + CO 2 + H 2 O, 

Marble Hydrochloric Calcium Carbon Water 
acid chloride dioxide 

we know that 2 X (1 + 35-5) grams of hydrochloric acid wil) 
give (12 + 32) grams of carbon dioxide, since Cl = 35-5 and 
C = 12. But (12 + 32) or 44 grams of carbon dioxide is the 
G.M.W. of this gas. Therefore at N.T.P. it will occupy the 
G.M.V. or 22-4 litres. Suppose we wished to know what 
weight of hydrochloric acid we must use in the above reaction 
to obtain 20 litres of dry carbon dioxide at 37 C. 700 mm. 
pressure. First find what volume this would be at N.T.P. 

20 X 273 X 700 _ c 00 v , 
310 X 760 = 16 ' 22 lltre8 ' 

Now 22-4 litres of carbon dioxide are obtained by using 
73 grams of hydrochloric acid, 

/. 16-22 litres are obtained by using 

73 X 16-22 

2^r = 52-9 grams. 

Calculation of a Formula from Percentage Corn- 
position. In determining the formula of a compound, one 
usually begins by analysing the compound quantitatively, 
the results of the analysis being expressed as a percentage 

1 Bear in mind that the molecular weight is simply a RATIO, whereas 
the gram-molecular-weight is a definite WEIGHT of the substance. 



composition ; thus the quantitative composition of calcium 
carbonate is 

Calcium 40 per cent., Carbon 12 per cent!, and Oxygen 48 
per cent, by weight. 

Does the percentage composition enable us to calculate the 
formula of a substance ? The percentage by weight of an ele- 
ment present in a given compound is obviously directly propor- 
tional to (a) the number of atoms of that element per molecule 
of compound, and (b) the atomic weight of the element. It is 
therefore proportional to the product of these two things, or 

Percentage of Element A in the compound C varies as atomic 
weight of A x number of atoms of A in 1 molecule of C. 
.'. No. of atoms of A in 1 molecule of C 

, ^ Percentage of A in C 

is proportional to : : - r r- 

Atomic weight of A 

Similarly, the number of atoms of the element B in 1 molecule 

- , , , ~ . , , Percentage of B in C 

of the compound C is proportional to : . T r-^r- 

Atomic weight of B 

and so on. 

.*. No. of atoms of element A in 1 molecule of compound C 
No. of atoms of element B 

Percentage of A in C Percentage of B in C 

Atomic weight of A * Atomic weight of B 
That is, the ratio of the numbers of atoms of the elements 
A, B, etc., which are present in the molecule of a compound C 
is the same as the ratio of the numbers obtained by dividing 
the percentages of A, B, etc., in the compound C by their respec- 
tive atomic weights. 

If, for instance, taking the above example, we divide the 
percentages of the elements calcium, carbon and oxygen in 
calcium carbonate by the atomic weights of these elements, 

we have 


% of calcium = 40 ; At. Wt. of calcium = 40 ; /( ? - == 1 

At. Wt). 


% of carbon =- 12 ; At. Wt. of carbon = 12 ; ./ = 1 

At. Wt. 


% of oxygen = 48 ; At. Wt. of oxygen = 16 ; = 3 

From these results we conclude that the ratio of the numbers 
of atoms of Calcium, Carbon and Oxygen present in a mole- 
cule of calcium carbonate is 1:1:3. 

Again, the percentage composition of alcohol is 

C = 52-2 per cent. ; H = 13-0 per cent. ; O = 34-8 per cent. 
Applying the same method, the ratio of the numbers of carbon, 
hydrogen and oxygen atoms present in the molecule of 

52-2 1 *VO ^4.8 

alcohol is \ : -^ : ~ = 4-35 : 13-0 : 2-17. Simplifying 
\.2i J. JLo 

this ratio by dividing each number by the highest common 
factor (in this case, 2-17 is the highest common factor), it 
becomes 2:6:1. Therefore the atoms of carbon, hydrogen 
and oxygen present in the molecule of alcohol are in the ratio 
of 2 of carbon to 6 of hydrogen to 1 of oxygen. 

This ratio is fulfilled in all the following formulae : 

C 2 H 6 ; C 4 H 12 O 2 ; C 6 H 18 O 3 ; etc. up to C^H^O,. 

In other words, the percentage composition alone of a substance 
is not sufficient to give us the true, formula of that substance. 
It gives us the simplest or empirical formula only, and the 
true formula is n times this, where n is a whole number since 
atoms are indivisible. Thus in the case of alcohol, the 
empirical formula is C 2 H 6 0, but the true formula is 
[C 2 H 6 O]n, and n we have still to discover. What further 
evidence is necessary to give us n, and therefore the true 
formula ? We can soon think of this by considering two or 
three possibilities for the true formula for alcohol, e.g. : 
(i) C 2 H 6 0; (ii) C 4 H 12 2 ; and (iii) C 6 H 18 3 . The mole- 
cular weight of (i) is 46, that of (ii) is 92, and that of (iii) is 
138. It is therefore possible to decide what multiple of the 
empirical formula of a substance is the true formula, by 
observing which gives a molecular weight corresponding to 
the observed value of the molecular weight. The vapour 
density of alcohol, calculated by Victor Meyer's method, is 
23 ; therefore its molecular weight = 46. Now a molecular 


weight of 46 corresponds to formula (i) above, and will not 
agree with any other ; hence here n = 1 and the empirical 
formula happens to be the true one also. This is not always 
the case. The empirical formula for grape -sugar is CH 2 0, but 
the molecular weight of this compound is found to be 180. 
The molecular weight of CH 2 = 12 + 2 + 16 = 30 ; there- 

1 80 
fore n = r^- = 6, and the true formula for grape-sugar is 

C 6 H 12 6 . 

It should be mentioned that we are ignorant of the mole- 
cular weights of many solid compounds, and in these cases 
we assume the true formula to be the same as the empirical 
formula until facts are discovered which show us that our 
assumption is untenable. Thus we write NaCl for solid 
common salt, but for all we know to the contrary its true 
formula might be Na 1000 Cl 1000 or Na 5 Cl 5 , etc. As we have no 
evidence either way, it is preferable to choose the simplest 
of these formulae. Modern work seems to show that, in such 
electrolytes (p. 94) as sodium chloride, etc., any single crystal, 
however large, may be regarded as a single molecule. For 
substances of this nature, the empirical formula is consequently 
the only one of practical value. 

The empirical formula for oxalic acid is CH0 2 , but oxalic 
acid will form salts in which only half the hydrogen is replaced 
by a metal. Therefore the true formula must be at least 
C 2 H 2 O 4 , and may be still more complex. Similarly the 
empirical formula of naphthalene is C 6 H 4 , but chlorine will 
replace one-eighth of the hydrogen in the naphthalene molecule. 
Hence the true formula for naphthalene must be at least C 10 H 8 
and may be still more complex. 

In short, the percentage composition of a substance enables 
us to calculate its empirical formula ; to get its true formula 
we need additional evidence, which is often directly obtainable 
by molecular weight determinations, but is sometimes of an 
indirect nature. 


1. What does a chemical equation tell you, and what does it not 
tell you, about the reaction which it represents ? 


2. Explain the terms gram-molecular-weight and gram-molecular- 

3. What volume of acetylene, C 2 H 2 , measured dry at 15 C. 740 mm., 
could be obtained by the action of water on 20 grams of calcium 
carbide, CaC 2 ? 

CaC a + 2H 2 O Ca(OH) 8 -f C a H 2 . 

4. What is the empirical formula of a substance ? How is it related 
to its true formula ? 

6. Cobaltite has the following percentage composition : 
Cobalt, 35-53 per cent. Arsenic, 45-18 per cent. Sulphur, 19-29 per cent. 

Find the empirical formula of the mineral. 

(Co = 59. As = 75. S = 32.) 

6. Olivine has the following percentage composition : 
Magnesium, 34-3 per cent. Silicon, 20-0 per cent. Oxygen, 45-7 per 

Find its empirical formula. 

(M = 24. Si = 28. O = 16.) 

7. A platinum compound has the following percentage composition : 

Platinum, 62-56 per cent. Nitrogen, 7-55 per cent. Hydrogen, 1-62 
per cent. Chlorine, 38-27 per cent. 

From these data find its empirical formula. 

(Pt = 195. N = 14. H = 1. Cl = 35-5.) 

8. An organic compound contains the following percentages : 
Carbon, 12-76 per cent. Hydrogen, 2-13 per cent. Bromine, 85-11 

per cent. 

Its vapour density is 94. Find its formula. 

(C =12. H - 1. Br = 80.) 

9. Sodium glycocholate has the following composition : 
Carbon, 64-07 per cent. Hydrogen, 8-62 per cent. Nitrogen, 2-87 per 

cent. Oxygen, 19-71 per cent. Sodium, 4-73 per cent. 
Find its empirical formula. 

(C - 12. H 1. N = 14. O - 1. Na = 23.) 


The behaviour of gases expressed by the gas laws, Avogadro's 
Hypothesis, the phenomena of evaporation and vapour- 
pressure, the fact that all gases are completely miscible with 
one another, and many other facts, are simply explained by 
the Kinetic Theory. This theory is largely due to CLAUSIUS 
(1857), BOLTZMANN (1868), and CLERK-MAXWELL (1860), and 
is shortly as follows : 

The molecules of a gas are in constant and rapid motion, 
and at ordinary pressures the spaces between the molecules 
are very large in comparison with the actual size of the 
molecules. If, for example, all the hydrogen molecules in a 
gas- jar " full " of the gas were to be packed so closely 
together that they were in actual contact, they would take 
up no more than a very minute fraction of the volume of the 
jar. How is it, then, that the jar appears to be full of hydro- 
gen ? Suppose you were in a room in one corner of which 
was a small box of golf-balls. The volume of the golf-balls 
would be negligible in comparison with the total volume of 
the room. But suppose that by some means those golf -bails 
began to move about in the room at the rate of 60 miles or so 
a minute ! In what part of the room would you feel safe ? 
The effective volume of the golf-balls, that is, the volume 
throughout which their effect is felt, would now be the whole 
room. Now, hydrogen molecules do move with an average 
velocity of rather more than 60 miles a minute ; hence a 
number of them that might occupy only a fraction of a cubic 
millimetre in actual volume would appear to fill a large flask 
if placed in it. 



It is easy now to see why gases will all mix with one another ; 
there are such comparatively wide spaces between their 
molecules and the latter move so quickly that mixing has 
nothing to prevent it and would in fact be expected. 

The pressure which a gas exerts is considered to be due to 
the bombardment of the walls of the vessel by the gas mole- 
cules. Think of one of those machines that invite you to 
" test your punch." If you hit it, a momentary pressure is 
recorded, and if you are quick enough you can hit again 
before the pointer has gone back to zero. If you could 
administer several thousand punches a second, all of about 
the same strength, the pointer would indicate a steady 
pressure. In a volume of gas confined in a closed space, the 
walls of the containing vessel get hit many millions of times 
in a second, and gas-pressure is caused in this way. This 
gives us an explanation of Boyle's Law. If we halve the 
volume in which a gas is confined we shall clearly double the 
number of hits which the molecules make on the walls of the 
vessel, that is, we shall double the pressure. Similarly, if 
the volume be increased to 3 times its original value, there 
will be only of the number of molecular impacts in a given 
time, and therefore the pressure will be reduced to of its 
original value. 

Suppose we have n molecules of mass m and velocity 
c enclosed in a cubical vessel with length of side I. If we can 
find the pressure on unit area of the walls of this vessel caused 
by a single molecule, and then multiply this pressure by the 
total number of molecules, we shall have the pressure per 
unit area caused by the gas as a whole. 

By the laws of symmetry we may assume, for purposes of 
calculation, that J of the total number of molecules are 
moving with uniform velocity c between one pair of parallel 
faces of the cube, another third between one of the remaining 
pairs of parallel faces, and the last third between the final 
pair of parallel faces. The time which any one molecule 
takes to travel from one face of the cube to the parallel face 

will be -, and therefore the number of hits per second that 


it makes on these two faces will be y, or -^ on one of the faces. 

i 1 % 

The molecule is moving with a velocity c and its mass is m, 
therefore its momentum is me, and if we suppose the mole- 
cules and the walls to be perfectly elastic, the molecule will 
rebound from the wall after impact with a momentum of me. 
Therefore the total change of momentum is me ( me ) = 2mc. 
This is balanced by the equal and outward reaction in the 
wall, that is, the momentum acquired by the wall from each 

impact is 2it?.c. Now there are -, impacts per second on one 

face from one molecule, therefore the total effect per second 
produced on one face by one molecule 
c _ me 2 

= a x 2wc = T- 

Hence the total effect produced on one of the faces by all the 

i ^ ^ t -n i n mc * 

molecules moving towards that face will be - X y-. 

5 i 

The total area of this face == I 2 , therefore the effect or 

., . ,, - , nmc 2 nmc 2 

pressure per unit area of the face = . j - = . =5- , 

i X ' v 

, nmc 2 . . A1 . . x , . 

i.e., p = J . - where v is the volume of the cube, or 

pv = J . nmc 2 . 

By a simple manipulation of this formula, 

2n me 2 


Now -^- is the kinetic energy of one molecule of the gas, and 

is the kinetic energy of all the molecules of the gas. 

It seems legitimate to assume that the kinetic energy of 
the molecules of a gas is a measure of the temperature of 

,. i o nmc 2 _ nmc 2 __ . _, 

the gas, whence f . ^ <x T, or J . ^ = RT, where R 

2 2i 

is a constant. And since the pressure of a gas is directly 
proportional to the kinetic energy of its molecules, that is, 


to their speed since their mass is constant, we arrive at the 
conclusion that the pressure exerted by a gas is directly 
proportional to the absolute temperature if the volume is 
constant a law with which we have met before (p. 30). 
Again, at absolute zero, when T = 0, RT = and 

| . also = 0, that is, the molecules will have no motion 
at all. 

Boyle's Law also follows from the equation pv = f . 


for \ . will be constant at constant temperature and 

therefore p must vary inversely as v. 

Charles' Law, too, is shown to have a theoretical basis, for 

pv == f . RT ; now if p is constant, v must vary 

directly as f . ^ which varies as T since R is constant. 

If we assume that equal volumes of gases at the same 
temperature and pressure have equal kinetic energies, which 
is more reasonable to assume than that they have different 

, . t . . A . 9 nmc 2 rim'c' 2 

kinetic energies, then f . ^ = f . ^ , where n and n 

are the respective numbers of the molecules in the equal 
volumes of the two gases, and m and m' their respective 
masses. Then since average kinetic energies of the molecules 
of these gases are assumed to be the same, 

me 2 _ m'c' 2 
_ =-g". 

mr z m'r"* 

But i.n. 7 ^- =I.n'. ^- 

f . n = * . n' 

or n = n'. 

In other words, equal volumes of gases at the same tempera- 
ture and pressure contain equal numbers of molecules, which 
is Avogadro's Hypothesis. 

. By further assumptions and measurements, it has been 



found possible to calculate the number of moleoules in any 
given volume of gas at a stated temperature or pressure. 
Approximately, 1 c.c. of a gas at N.T.P. contains 2-70 X 10 19 
molecules. The G.M.V. of a gas at N.T.P. contains about 
6-06 x 10 23 molecules ; this number is called Avogadro's 

Velocity of Molecules. In the equation pv = f . , 


suppose that the volume is the G.M.V. (22-4 litres) and the 
pressure 760 mm. and that the gas is at C. Then the 
kinetic energy of this volume of gas will be the kinetic energy 
of a gram-molecule of the gas, since the G.M.V. is the volume 
of the gram-molecule of any gas at N.T.P. 
.*. The kinetic energy at N.T.P. of the gram-molecule of a 

nmc 2 
gas, i.e., -y-, 

B | pv = | x (760 X 13-6 x 981) X 22,400 ergs. 
= about 34 x 10 9 ergs. 

This is called the molecular energy of translation of a gas at 
N.T.P., and from it we can calculate the speed of the molecules, 
c ; thus, at N.T.P. 


34 X 10 9 = -^- where nm M, 1 gram molecule. 


In the case of hydrogen M = 2, 

, 2 x 34 x 10 9 
' C = - M - ' 
/. c 2 = 34 X 10 9 .'. c = V34 X 10 9 
= about 170,000 cm. per sec. 

For carbon dioxide, c 2 = 

2 X 34 X 10 9 . /34 X 10 

/. c = V - 

44 ' ' v 22 

= about 36,200 cm. per sec. 1 

1 These velocities are only approximate average velocities ; the actual 
velocities of the molecules of a specimen of gas vary considerably 
from molecule to molecule. 


Rate of Diffusion of a Gas (Graham's Law). GRAHAM 
found that the rates at which two gases passed through a 
porous partition were inversely proportional to the square 
roots of the densities of the gases, or 

R' VD 
R" * VW'' 

[This phenomenon of the passage of a gas through a porous 
partition is more properly called effusion, " diffusion " being 
reserved to describe the passage of one gas (or liquid) into 
another when they are not separated by a partition.] 

Graham's Law may be deduced from the equations of the 
kinetic theory. For 

me 2 m'c'* c 2 m' e 

2 2 '" c' 2 m " c' m 

that is, the velocities of the molecules of two gases at the 
same temperature and pressure are inversely proportional to 
the square roots of their masses, and therefore to the square- 
roots of their densities. 

An apparatus, the Effusiometer, was devised in 1858 by 
BUNSBN to measure the rates of effusion of gases, and in this 
way to compare their molecular weights. The apparatus 
consists of a tube inverted in a trough of mercury and closed 
at the upper end by a platinum plate in which is a very fine 
hole. The gas under observation is placed in the tube, which 
is then sunk in the mercury to such a depth that the gas is 
under a suitable pressure, when it slowly escapes through the 
fine hole in the platinum. By measuring the volumes of two 
gases that escape through the hole in a constant time the 
ratio of the two molecular weights may be calculated. 

Evaporation and Vapour Pressure. The molecules of 
a liquid, like those of a gas, are in motion, but they do not 
move so fast as those of gases. However, if a surface of 
liquid be exposed, .some of the molecules which come to the 
surface will have sufficiently high velocities to leave the 
liquid altogether. Continuation of this process will finally 
cause the entire conversion of the liquid into gas, if the 


molecules which leave the liquid never return to it. If, 
however, the liquid be placed under a bell- jar, the escaping 
molecules have only a limited space in which to move, and 
after a time some of them will strike the liquid again and will 
be absorbed in it. When the number of molecules that 
leave the liquid in a given time is equal to the number that 
return to it in that time, no apparent change will take place, 
and the space in the bell- jar will be saturated with the vapour. 

A current of air blown over the surface of a liquid will sweep 
away all the molecules that come off from it and will there- 
fore aid evaporation. 

The cause of vapour pressure will be apparent from the 
above considerations. 

Deviations from the Gas -Laws. No gas obeys the gas- 
laws exactly. The deviations that occur are considered to 
be due to two things, (i) the actual volume of the molecules 
of the gas, and (ii) to their mutual attraction. For a given 
mass of an " ideal " gas at constant temperature, PV should 
be absolutely constant ; in practice it is found that PV is 
only approximately constant, the deviations becoming greater 
as the pressure increases. Let us see how (i) and (ii) above 
will affect the behaviour of a gas. Under ordinary pressures, 
the actual volume of the molecules will be negligible in com- 
parison with their effective volume, and the molecules will be 
so far apart from one another that any attraction they have 
for one another will also be so small as to be negligible. Hence 
under these conditions PV will be very nearly constant ; it 
may in fact be so nearly constant that deviations cannot be 

But suppose that the pressure is increased so much, and 
the volume proportionately diminished, that the actual 
volume of the molecules is an appreciable fraction of the 
volume in which they are confined. It is clear that the 
molecules will strike the walls more often than they would 
if they were mere points. Thus, if the diameter of a molecule 
is 1 unit and the distance between two parallel faces of the 
vessel in which it is confined 5 units, instead of having to 
travel 5 units from one wall before it strikes the other, it will 


have to travel only 4 units, and the " pressure " it exerts will 
therefore be increased proportionately. If, on the other 
hand, the distance between the walls is 100,000 units, it will 
have to travel 99,999 units ; now the difference between 
100,000 and 99,999 is probably too small to be detected in 
ordinary experiments, whereas the difference between 5 and 4 
is 20 per cent, of the former number and therefore easily 

In short, the fact that molecules have a definite volume 
and are not mere points causes the pressure to increase too 
rapidly for PV to remain constant ; or, what comes to the 
same thing, the volume V in the expression PV should repre- 
sent the volume in which the molecules move. This will be the 
volume occupied by the gas minus an incompressible part 
which is proportional to the actual volume of the molecules. 
VAN DER WAALS allowed for this by writing P(V 6) = a 
constant, where 6 is a constant proportional to the actual 
volume of the molecules. 1 This constant, 6, is known as the 
molecular co-volume. 

Turning now to the second factor in the deviation from the 
gas-laws, namely, the attraction of the molecules for one 
another, it is evident that this will tend to draw the molecules 
closer together, and therefore to make the volume decrease 
too rapidly. Van der Waals corrected for this by writing 

( P + ^2 ) instead of P, where a is a constant and ^ is a 

measure of the molecular attraction. 
Van der Waals' Equation therefore is 

+ ^ (V - b)= RT 

and gives results that correspond fairly well with the experi- 
mental figures. A more exact equation was suggested by 
DIETEBICI in 1899, but for ordinary pressures van der Waals' 
equation is sufficiently accurate. 

Critical Phenomena. When gases are cooled to a 
sufficiently low temperature, they become liquids, even under 
1 6 is not the actual volume of the molecules. 


atmospheric pressure. If the pressure is increased, lique- 
faction may be brought about at higher temperatures, but 
for each gas there is a certain temperature above which it is 
impossible to liquefy that gas, whatever may be the pressure 
applied. This temperature is called the critical temperature 
of the gas, and the minimum pressure which is required to 
bring about liquefaction at the critical temperature is called 
the critical pressure. The first chemist to study critical 
phenomena seriously was ANDREWS, who in 1869 described 
some of his results in these words : " On partially liquefying 
carbon dioxide by pressure alone, and gradually raising at 
the same time the temperature to 31-1 C., the surface of 
demarcation between the liquid and gas became fainter, lost 
its curvature, and at last disappeared. The space was then 
occupied by a homogeneous fluid, which exhibited, when the 
pressure was suddenly diminished, or the temperature slightly 
lowered, a peculiar appearance of moving or flickering striae 
throughout the entire mass. At temperatures above 31-1 C. 
no apparent liquefaction of carbonic acid, or separation into 
two distinct forms of matter, could be effected, even when a 
pressure of 300 or 400 atmospheres was applied." Other 
gases gave similar results. 

The explanation of the phenomenon is that at the critical 

temperature and pressure the specific volumes, i.e., -r r-~, 

of the gas and liquid become the same. 

Liquefaction of Gases. All known gases have been 
liquefied. Helium gave the most trouble, but in 1907 ONNES 
succeeded in preparing liquid and even solid helium. Many 
gases can be liquefied at C. by application of pressure only ; 
in other words, their critical temperatures lie above C. 
Such are sulphur dioxide, carbon dioxide, ammonia, chlorine 
and cyanogen (C 2 N 2 ). Others, such as oxygen, hydrogen, 
methane (CH 4 ) and nitrogen have first to be strongly cooled. 
On account of the apparent impossibility of liquefying these 
gases, they were formerly called the " permanent " gases ; 
but when it was realized that a low temperature and not a 
high pressure was the important factor it was soon found 



possible to liquefy them. The necessary low temperatures 
were at first obtained by allowing liquid carbon dioxide to boil 
rapidly, when the absorption of the latent heat of vaporization 
causes the temperature of the remaining liquid to fall lower 
and lower. This method has, however, been superseded by 
another process, based on 
a different principle called 
the Joule-Thomson effect, 
discovered by JOULE and 
later became LORD KELVIN). 
When a gas is compressed 
and then allowed to escape 
through a porous plug into 
a region of low pressure, it 
becomes cooler. This de- 
pends upon the fact that in 
the compressed gas the 
molecules are close together, 
while after passage through 
the plug they are much 
farther apart, and therefore 
the force of attraction that 
they have for one another 
must have been overcome. 
To overcome this attraction 
energy is required, and this 
is supplied by the heat of 
the gas itself, which there- 
lore becomes cooler. 

If the cold gas thus pro- 
duced is allowed to circulate 

around the tube that contains the compressed gas, the latter 
will be cooled before it reaches the plug, and will therefore 
be colder still after passage through the plug. By continu- 
ing this process, the gas will get colder and colder and will 
finally issue from the plug as a liquid (Fig. 14). 

The liquefaction of gases is now a regular industry, and for 

14. Liquefaction Apparatus. 
A. Valve screw. B. Receiver. 


details of the various processes employed the technical books 
should be consulted. 

Liquid air, etc., can be preserved for some days in DEWAB 
or vacuum vessels, which are double-walled vessels with a 
high vacuum between the walls. The inner surfaces of the 
walls are silvered in order to reduce absorption of heat, while 
the vacuum allows very little heat to pass through it. 

FIG. 15. Dewar Flasks. 


1. What is the kinetic theory of gasos ? 

2. How are evaporation and vapour pressure explained by the 
kinetic theory ? 

3. Show how Boyle's Law may be deduced theoretically. 

4. Explain the terms critical pressure and critical temperature. 

6. What are the main principles involved in the various methods 
employed for the liquefaction of gases ? 

6. Show that Avogadro's Hypothesis has a theoretical basis, if the 
kinetic theory o f gases is " true.'' 


The formula for ammonium chloride is NH 4 C1 ; its mole- 
cular weight is therefore 53-5, and we should expect its- 
vapour density to be 26-75. Experiments show, however, 
that its vapour density is 13-37. How are we to explain 
this, and what does it mean ? Let us consider the possi- 

(i) If our previous work is correct, then the molecular 
weight of ammonium chloride should be twice its vapour 
density, that is 26-75. But one molecule of ammonium 
chloride cannot contain less than one atom of chlorine, and 
the atomic weight of chlorine is 35-5 ! We could assume that 
the molecular weights of nitrogen and chlorine had been in- 
correctly established, and that instead of being 14 and 35-5 
they were 7 and 17-75 respectively. This would give us a 
formula NH 2 C1 for ammonium chloride, for N now = 7, 
Cl = 17-75, and H = 1 /. NH 2 C1 -= 26-75, which corresponds 
with the vapour density of 13-37. 

Unfortunately, however, this explanation at once lands us 
in far more serious difficulties concerning the other compounds 
in which nitrogen and chlorine are found. Thus, the mole- 
cular weight of hydrochloric acid calculated from its vapour 
density is 36-5. Now if the atomic weight of chlorine is 
17-75, the formula for hydrochloric acid will be HC1 2 , where 
the chlorine has a, valency of 0-5 a fact that at once arouses 
our suspicions. Moreover, the ratio of the specific heats of 
hydrochloric acid gas agrees with that of a diatomic gas and 
differs from that of a triatomic. Hence we must abandon 



the idea that the atomic weights of chlorine and nitrogen 
are half the generally accepted values. 

(ii) Still assuming that the molecular weight of ammonium 
chloride vapour is twice its vapour density, we might argue 
that the formula for ammonium chloride is NjH 2 Clj. But 
this involves a splitting of atoms and therefore an abandon- 
ment of the atomic theory. We should, therefore, not accept 
this suggestion unless no alternative could be found. 

(iii) We might suppose that Avogadro's Hypothesis was 
not universally true, and that ammonium chloride vapour 
was one of the cases in which the Hypothesis was inapplicable. 
But, as we have seen, the Hypothesis has a theoretical basis, 
and chemists would be reluctant to admit the possibility of 
exceptions to it. We should, indeed, be explaining a difficulty 
by creating a greater one. 

(iv) The suggestion finally adopted by chemists was that 
the ammonium chloride vapour really does not consist of 
ammonium chloride molecules, but of a mixture of ammonia 
and hydrochloric acid molecules in equal numbers 
NH 4 C1 = NH 3 + HC1. 

A mixture of these two gases in equal volumes will obviously 
have a density half-way between that of ammonia, 8-5, and 

8'5 -4- 18*25 
that of hydrochloric acid, 18-25, that is ^~ = 13-37 


which is the observed value. 

When the vapour is cooled, ammonium chloride is re- 
formed. It is evident that this theory is much more satis- 
factory than any of the others, especially as it is quite easy 
to prove that some ammonia and hydrochloric acid are present 
in the vaporized ammonium chloride, although it is not easy 
to demonstrate the complete splitting-up. 

Unexpected values for the vapour densities of many other 
substances have been obtained, and in every case it has been 
found possible to explain the observed values by assuming 
that a splitting-up of the original molecules has occurred. 
We can define the phenomenon as follows : 

When the molecules of a substance split up, on heating, into 
simpler molecules that recombine on cooling, the phenomenon 


is called DISSOCIATION, and the substance is said to 
dissociate. 1 

Other substances that dissociate on heating are 
iodine (I 2 - ^ I + I), 

nitrogen peroxide (N 2 O 4 ; ^ N0 2 + N0 2 ), 
phosphorus pentachloride (PC1 5 N PC1 3 + C1 2 ), 
hydrogen iodide (2HI - ^ H 2 + I 2 ; ^ H 2 + I + I), 
calcium carbonate (CaCO 3 N CaO + CO 2 ), 
and many more. 

These actions are, then, reversible, the direction in which 
they proceed depending upon the external conditions of 
temperature, and, in most cases, pressure. To indicate this 
reversibility the sign s ^ is used instead of == . 

Nitrogen Peroxide. Nitrogen peroxide is a brown liquid 
that boils at 22 C. At temperatures just above the boiling- 
point its vapour density is 46, but as the temperature rises 
the vapour density decreases until it reaches a minimum 
value of 23 at 140 C., as shown in the following table : 

Vapour Vapour 

Temperature. Density. Temperature. Density. 

22-5 .... 45-8 100 .... 24-3 

27 . . . . 38-3 135 .... 23-1 

60 .... 30-2 140 .... 23-0 

These figures may be explained by assuming that at 22-5 the 
vapour of nitrogen peroxide consists almost entirely of N 2 04 
molecules, having a molecular weight of 92 and therefore a 
vapour density of 46, which corresponds to the observed value. 
On heating, some of the N 2 O 4 molecules split up each into 
two NO 2 molecules, until at 140 C. they have all split up and 
the gas now consists entirely of N0 2 molecules of molecular 
weight 46 and vapour density 23. At intermediate tempera- 
tures, then, the gas will consist of a mixture of N 2 O 4 molecules 
with NO 2 molecules. The percentage of N0 2 by volume or 
by weight at any given temperature may be calculated as 
follows from the observed vapour density, n : 

1 If the splitting-up of the substance is irreversible, it is called 
decomposition and not dissociation. 


(a) Per cent. N0 t by volume. 

Let x = % NO 2 by volume, when the V.D. of the mixture is n. 
A 100 - x = % N 2 4 ^ 

The density of NO 2 is 23 ; that of N a O 4 is 46. 


Now density = . 

.'. weight of the NO a == x X 23 
and weight of the N 2 O 4 = (100 s)46 

.'. total weight of gas = 23z + 46(100 x). 

But there are 100 parts by volume of the gas and its density is n, 
.*. total weight of the gas = 100 X n. 
:. 23x + 46(100 - a?) = 100 X n. 
Hence for any given value of n, x may be calculated. 

(b) Per cent. NO% by weight. 

Let x % NO 2 by weight, when V.D. of the gas is n. 

:. 100 - x % N 2 o 4 

V.D. of N0 2 = 23 ; that of N 2 O 4 is 46. 

^ . weight 

Density = ^ - 


/. volume of the NO 2 = -~ 


and volume of the N 2 O 4 = 

.*. total volume of the gas = - -\ . 

2t& 4u 

But there are 100 parts by weight of the gas and its density = ti 
.'. total volume of the gas = 

x 100 x __ 100 
" 23 ~* 46 ~" ~n~' 
Hence for any given value of n, x may be calculated. 

Effect of Changes of Pressure upon Dissociation. 

This will be discussed again later (pp. 123 j^.), but we can con- 
sider it shortly here. When nitrogen peroxide, N 2 4 , dis- 
sociates, according to the equation N 2 O 4 % N . 2NO 2 , it is 
seen that 2 molecules are formed from 1 original molecule. 
If therefore the pressure remains constant, the volume of the 
gas will be doubled, by Avogadro's Hypothesis. Suppose 
now we take this volume of NO 2 molecules and increase the 
pressure, keeping the temperature constant. What will 
happen ? The volume having decreased, the N0 a molecules 


will have less room in which to move about and will therefore 
meet one another more often, and the chances of two NO 2 
molecules reeombining to form an N 2 O 4 molecule will be 
increased. In other words, in this ease increase of pressure, 
will cause a decrease of dissociation. This is true in every case 
in which the products of dissociation occupy a larger volume 
than the original gas under the same conditions of temperature 
and pressure. 

We shall find that the dissociation of hydriodic acid gas, 
however, is not accompanied by an increase in volume, for 

2 HI ^= H 2 + I 2 , 

i.e. we have started with 2 molecules and get 2 molecules 
after dissociation, 1 therefore there is no change in volume. 
In this case we should expect change of pressure to have no 
effect upon the degree or extent of dissociation, and this is 
confirmed experimentally. 

Le Chatelier's Principle. We can consider the effect 
of pressure (and also of temperature) on dissociation in terms 
of a principle enunciated by the French chemist LK CHATELIER. 
This is expressed as follows : - 

// to a system in equilibrium a constraint be applied, a 
change takes place within the system tending to nullify the 
effect of the constraint and to restore the equilibrium. 

This is a general law of chemistry, and we can understand 
very well what it means by an analogy from everyday life. 
Suppose a man to have regulated his expenditure in such a 
way that his expenditure + income-tax = total income. 

We have here a '' system in equilibrium." Let us now sup- 
pose that the Government apply a " constraint " by increasing 
the rate of income-tax. A change will have to take place 
within the system tending to restore the equilibrium which 
has been upset. If the income is fixed, the expenditure will 
have to be lowered and this is an example of Lo Chatelier's 
principle working, in everyday life ! 

Let us take now a chemical example. Suppose we have 

1 Assuming that the temperature is sufficiently high Jor the iodine to 
remain as vapour. 



some water in a cylinder fitted with a piston (Fig. 16). The 
space above the water is saturated with water-vapour, and 
the temperature is constant. 

If as many H 2 O molecules leave the water for the vapour 
in a given time as return from the vapour to the water in 
that time, the system will be in equilibrium. If now 
the piston is pushed down a little way, the pressure on 
the vapour will have been increased, and therefore by Le 
Ch atelier we should expect a change to take place within 
the system the result of which would be to nullify the 
effect of pushing down the piston. In point of fact, what 
happens is that some of the water- vapour is condensed to 

water vapour 

HEF ^ wafer 

FIG. 16. 

FIG. 17. 

water ; for since the volume is decreased, more molecules 
will pass from the vapour into the water in a given time 
than will leave the water in that time. When sufficient 
vapour has condensed in this way, the numbers of molecules 
leaving and entering the water per second will become equal 
again, and equilibrium will be restored. 

For a second example, take ice at and water at in 
a similar cylinder (Fig. 17). If no heat enters or leaves the 
cylinder, the ice and water will remain unchanged, that is, 
the system will be in equilibrium. Now suppose we put 
weights on the piston we shall be applying a " constraint " to 
the system, and according to Le Chatelier a change should 
take place within the system tending to nullify the constraint, 
i.e. to reduce the pressure. This can be done only by a 


reduction of volume. Since 1 gram of ice at occupies a 
greater volume than 1 gram of water at 0, a reduction of 
volume could occur if the ice were to melt, and this in fact 

Applying now Le Chatelier's principle to cases of dissocia- 
tion, it follows that if dissociation is accompanied by increase 
in volume, an increase in pressure on such a system in 
equilibrium will tend to reduce the dissociation, since in this 
way the volume and therefore the pressure will be lowered. 
If, however, dissociation is accompanied by no change of 
volume, change of pressure will have no effect on the degree of 
dissociation, since the latter will have no effect on the volume. 

We have so far considered the effect of changes of pressure 
only, but by means of Le Chatelier's principle we can also 
predict the effect of temperature changes. Suppose the 

A^=B + C 

is accompanied by an absorption of heat in the forward 
direction, ->. What effect upon the degree of dissociation 
will heating produce ? Le Chatelier tells us that a change 
should take place tending to nullify the effect of heating, or, 
in other words, a change should take place which is accom- 
panied by an absorption of heat. This means that the dis- 
sociation will be increased. Similarly, if dissociation is 
accompanied by evolution of heat, heating will decrease the 


Chemical change is usually accompanied by an evolution 
or an absorption of heat. A reaction in which heat is evolved 
is said to be exothermic, while one in which absorption of heat 
takes place is called endothermic. Exothermic reactions are 
more common than endothermic. 

The heat evolved or absorbed in the formation of a com- 
pound from its elements is called the heat of formation of the 
compound. It is usually positive, 1 in which case the compound 
is exothermic, but is sometimes negative, 2 when the compound 
is said to be endothermic. 

l That is, heat is evolved. *That is, heat is absorbed. 



The number of calories evolved or absorbed in a particular 
reaction may be measured by carrying out the reaction in a 
calorimeter surrounded by a known weight of water and 
measuring the change in temperature of the water. A simple 
type of calorimeter for this purpose is shown in Fig. 18. For 
measuring the heats of combustion of substances in oxygen 
special calorimeters capable of withstanding a high pressure 

Fio. 18. 

have been constructed ; they are usually called " bomb- 
calorimeters. " 

The heat change of a reaction is expressed by adding the 
appropriate figures to the ordinary chemical equation repre- 
senting the reaction ; thus 

C + 2 == CO 2 + 97,000 calories 
means that when 12 grams of carbon combine with 32 grams 


of oxygen, 44 grams of carbon dioxide are formed, with 
evolution of 97,000 calories. Carbon dioxide is therefore an 
exothermic substance. 

C + 8 2 = CS 2 28,700 calories 

shows us that when 76 grams of carbon disulphide are pre- 
pared by direct combination of carbon and sulphur, 28,700 
calories are absorbed. Carbon disulphide is therefore an 
endothermic compound. 

A little thought will show us that to make these thermo- 
chemical equations quite definite, we ought to specify the 
physical state of all the substances represented in them. 
Thus the equation 

H 2 + oxygen = H 2 + 58,700 calories 

is true only if the hydrogen, oxygen and water are all in the 
gaseous state. If the water is allowed to condense, it is clear 
that the total heat evolved will be increased by the latent 
heat of vaporization given out by 18 grams of steam 
condensing to water, i.e. 9,700 calories ; hence 

H a + oxygen = H 2 + 58,700 calories + 9,700 calories. 

Gas. Gas. Liquid. 

Law of the Conservation of Energy. In no case 
hitherto investigated has any creation or destruction of 
energy been observed ; conversion of one form of energy, 
such as electricity, into another, such as heat, is a common 
phenomenon, but there is no loss or gain of energy in the 
process. This fact of experience is expressed in the law of 
the conservation of energy : Energy can neither be created nor 
destroyed. It is true that recent work has demonstrated that 
energy may be created by the annihilation of matter, and that 
the laws of the conservation of matter or energy ought really 
to be combined in a single law ; but such considerations need 
not delay us at- the moment. 

From the principle of the conservation of energy we may 
deduce the Law of Constant Heat Summation, sometimes 
known as Hess's Law, after its discoverer : The total amount 


of heat absorbed or evolved in a given reaction is independent of 
the number of stages in which the reaction is. brought about. 
This law enables us to calculate the heats of formation of 
many compounds that cannot be made directly from their 

Thus (i) we may calculate the heat of formation of carbon monoxide 
from the following data : 

C -f O 2 = CO 2 -f 97,000 calories. 
Solid. Gas. Gas. 

CO 4- oxygen = CO a -f 68,000 calories. 
Gas. Gas. Gas. 
Therefore, by subtraction, 

C 4- oxygen = CO -f (97,000 68,000) calories. 
Solid. Gas. 

or = CO + 29,000 calories. 

ii. Consider, again, the heat of formation of acetylene, C 2 H S . We 
know from experiment that 

C -f O a = CO a 4- 97,000 calories 
and H a -f- oxygen = H 2 O + 68,400 calories. 


The heat of combustion of acetylene is found experimentally to be 
310,200 calories 

C 2 H 2 + 2JO 2 - 2CO 2 4- H 2 O -f 310,200 calories. 

But the heat of formation of 2CO 2 4- H. 2 O will be 

2 X 97,000 4 68,400 = 262,400 calories. 

Hence the decomposition of 1 gram molecule of acetylene must be 
accompanied by the evolution of 310,200 262,400 calories, or the 
heat of formation of acetylene is 47,800 calories. 

iii. The heat of combustion of 1 gram molecule of alcohol, C 2 H 6 OH, 
is 340,500 calories, i.e., 

C 2 H 6 O 4- 3O 2 = 2CO 2 4- 3H 2 O 4- 340,500 calories. 
Liquid. Gas. Gas. Liquid. 
But the heat of formation of 2CO 2 and 3H 2 O will be 

2 X 97,000 4- 3 X 68,400 = 399,200 calories. 

Hence the heat of formation of alcohol is 399,200 340,500 calories, 
= 58,700 calories. 

Heat of Solution. When a substance is dissolved in water, 
there is usually either an evolution or absorption of heat. 
The heat change that occurs when 1 gram molecule of a 
substance is dissolved in a large quantity of water is called the 
heat of solution of that substance. In the thermochemical 



equation the water in which the substance is dissolved is 
written Aq. Thus 

NaOH + Aq = NaOH . Aq + 9,800 calories 

means that when 40 grams of caustic soda are dissolved in a 
large quantity of water, to form caustic soda solution 
(NaOH . Aq), 9,800 calories are evolved. 

Atomic Heats. The law of DULONG and PETIT has been 
mentioned previously. It states that the atomic weight of 
solid element X its specific heat is approximately equal to 
6*4, or the atomic heat of all solid elements is the same. At 
ordinary temperatures this is roughly true for metals, but 






FIG. 19. Curve showing Atomic Heats of Aluminium at various 

not for most non-metals. At higher temperatures the law is 
obeyed more exactly, while at the temperature of liquid 
hydrogen it does not hold at all. The atomic heats of all 
elements decrease as the temperature is lowered, a phenomenon 
which, with very many others, has been satisfactorily 
explained by PJanck's quantum theory. This is too 
advanced a subject for discussion here, but may be stated in 
the following words : Atoms emit or absorb energy in small 
definite amounts, directly proportional to the frequency of 


vibration of the atom ; they do not emit or absorb energy 
continuously. The " small definite amounts " of energy are 
called quanta. 


1. Write a general account of Thermal Dissociation. 

2. The vapour density of nitrogen peroxide at 60 is 30-2. Cal- 
culate the percentage by weight of NO 2 molecules present. 

3. State Le Chatelier's principle and explain its application, 

4. What is Hess's Law ? Show how it may be applied in the deter- 
mination of heats of formation. 

6. Calculate the heat of formation of methane, CH 4 , from the follow- 
ing data 

C -f O 2 = CO 2 + 97,000 calories. 
H a + |O a = H 2 O -f 68,400 calories. 
CH 4 -f 2O a = CO a -f 2H 2 O -f 211,800 calorie* 


NICHOLSON and CABLISLB showed in 1800 that if the wires 
from an electric battery are immersed in acidulated water 
the water is decomposed into hydrogen and oxygen. The 
hydrogen comes off from the wire attached to the negative 
pole of the battery, and the oxygen from that attached to 
the positive pole. This decomposition of a substance by the 
passage of electricity through it is called electrolysis. Pure 
water will not conduct electricity, 1 but if a little sulphuric 
acid be added conduction takes place and the water is 

The two metallic plates or wires that are placed in the 
liquid to bring about electrolysis are called the electrodes. 
The electrode connected to the positive pole of the battery 
is called the anode or positive electrode, and that connected 
to the negative pole is called the cathode or negative electrode. 
BEBZELIUS showed that aqueous solutions of acids, bases 
and most metallic salts will conduct electricity and are de- 
composed in the process. Thus, if a current of electricity 
is passed through a solution of copper chloride, chlorine is 
evolved at the anode (if this is made of a substance, like 
carbon, not attacked by chlorine) and copper is deposited 
on the cathode. Substances that are liberated at the anode 
are said to be electro-negative, since they are attracted to the 
positive pole, and it is a general rule of electricity that a 
positively-charged body repels another positively-charged 

1 This statement is not strictly accurate. Even the purest water 
will conduct the current to a slight extent. See p. 137. 




body but attracts a negatively- charged one, and vice versd. 
In the same way, substances that make their appearance 
at the cathode are called electro-positive. 

The phenomena of electrolysis were first carefully studied 
by FARADAY in 1832. He called solutions that are decom- 
posed when electricity passes through them electrolytes, but 
this name is now usually given to the substances dissolved 
and not to the solutions themselves. Non-electrolytes, then, 
are substances whose [aqueous] solutions do not conduct 

The substances that appear at the electrodes during 
electrolysis Faraday called ions (from the Greek, because they 
wander to the electrodes). Those that go to the anode 
are called anions, and those that go to the cathode cations. 
(Note that the modern significance of these terms is rather 
different from that used by Faraday. See Chapter XIII.) 
Diagrammatically we can represent these ideas as follows : 


Eleclrolylic Cell: 

FIG. 20. 

The conventional sign for a single cell of a battery is 
where the short, thick stroke represents the negative pole 





and the long, thin one the positive pole. A battery of 5 
such cells connected in series is shown in Fig 2 la, or more 
simply in Fig. 216. 

FIG. 21. 

If the positive pole of one cell is connected to the negative 
of the next, and the positive of this to the negative of the 
third, and so on, as shown in the diagram, the battery is said 
to be composed of cells connected in series. The terminal poles 
of the battery are then the negative pole of the cell at one 
end of the series and the positive pole of the cell at the other 
end. Fig. 22 shows 3 electrolytic cells and a battery con- 
nected in series. 


Cell 2 
FIG. 22. 

Cell 3 

Note that the electrode in Cell 1 connected with the negative 
terminal pole of the battery is also negative it is really an 
extension of the pole ; similarly with the anode in Cell 3. 

Faraday's Laws of Electrolysis. Faraday showed 
that the same quantity of electricity passed through solutions 
of electrolytes connected up in series will liberate in each cell 
weights of the products of electrolysis that are in the ratio of 
their chemical equivalents. Thus the same quantity of elec- 


tricity that liberates 1 gram of hydrogen will liberate 35-5 
grams of chlorine or 108 grams of silver or 31-7 grams of 
copper or 8 grams of oxygen, etc. 

The unit of quantity of electricity is the coulomb, which is 
the quantity of electricity conveyed by a current of 1 ampere 
flowing for 1 second. The number of coulombs required to 
liberate the equivalent in grams of any substance is found 
to be 96,000. A smaller or a greater number will liberate 
a proportionately smaller or greater weight. Faraday's Laws 
may therefore be expressed in the following terms : 

First Law. The weight of an ion liberated in electrolysis is 
proportional to the quantity of electricity which has passed 
through the electrolyte (i.e. to current X time). 

Second Law. The quantity of electricity required to deposit 
the equivalent in grams of an ion is 96,000 coulombs (or 1 

Theories of electrolysis will be dealt with later (Chapter 
XIII). For the apparatus used in various electrolytes, see 
under water, caustic soda, hydrochloric acid, etc. The 
quantity of electricity which flows through the circuit is 
measured by noting the time in seconds and multiplying this 
by the number of amperes of the current as shown by an 
ampere- meter or ammeter connected up in series in the circuit. 
An electrolytic cell is often called a voltameter ; this must not 
be confused with a voltmeter which is an instrument used 
for measuring the voltage or potential difference between two 
points of a conductor. 

It is also necessary to note that the primary products of 
electrolysis will often act upon water or upon one another, so 
that the products actually obtained may be secondary. Thus, 
when a solution of caustic soda is electrolysed, the primary 
products are sodium at the cathode and " hydroxyl " groups 
(OH) at the anode. But the sodium formed immediately 
acts upon the water so that the actual product at the cathode 
is hydrogen 1 

2Na + 2H 2 O = 2NaOH + H 2 , 

1 It is possible that the hydrogen may be the primary product 
after all, but I have gi\enr tla& usual explanation. 



while two hydroxyl groups react together at the anode, 
forming oxygen and water 

20H = H 2 + oxygen. 

Hence the final result of electrolysing a solution of caustic 
soda is that hydrogen is liberated at the cathode and oxygen 
at the anode, while the concentration of caustic soda round 
the cathode increases. In other words, it appears as though 
the water has been electrolysed, the weight of caustic soda 
remaining constant. The same sort of thing happens in the 
case of dilute sulphuric acid. Here the primary products of 
electrolysis are hydrogen and the " sulphate " group of atoms, 
S0 4 . The latter groups, however, when liberated at the 
anode, act upon the water present to give sulphuric acid 
and oxygen 

2S0 4 + 2H 2 - 2H 2 S0 4 + 2 , 

hence the actual product at the anode is secondary, namely, 

To take a rather different case, suppose we electrolyse a 
solution of copper sulphate, using a platinum or copper 
cathode and a copper anode. Here the copper particles go 
to the cathode, where they are deposited, while the S0 4 
groups at the anode dissolve the latter to form more copper 

Cu + S0 4 = CuS0 4 . 

The net result of electrolysis in this case is therefore the 
transference of copper from the anode to the cathode. This 
process is used in copper-plating, and similar ones in gold- and 
silver-plating, hence the term " electro-plate." 


1. Explain the terms electrolyte, anode, cathode, cation, anion. 

2. State Faraday's Laws of Electrolysis. 

3. The same quantity of electricity that deposits 3-53 grams of 
a metal in an electrolytic cell will liberate 453 c.c. of hydrogen, col- 
lected over water at 18 C. 763 mm. What is the equivalent of the 
metal f 


When a substance is dissolved in a liquid, the liquid is called 
the solvent , the substance dissolved the solute , and the product 
a solution. These terms are convenient, but it is easily seen 
that there may be cases in which they are rather difficult to 
apply systematically. For instance, if we dissolve 1 gram of 
alcohol in 100 c.c. of water, the alcohol is the solute and the 
water the solvent. Similarly, if we dissolve 1 gram of water 
in 100 c.c. of alcohol, here the alcohol is the solvent and the 
water the solute. But if we dissolve 50 c.c. of alcohol in 
50 c.c. of water, both substances have an equal right to be 
called solvent, or solute. In practice, however, there is no 
difficulty over such matters ; it is necessary merely to remem- 
ber that there is no essential property that marks off a 
solute from a solvent. It is just as correct to call brine a 
solution of water in salt as to call it a solution of salt in water 
it is, in fact, a homogeneous mixture of two things. Perhaps, 
therefore, we may define a solution as a homogeneous mixture 
of two or more substances. By this definition we see that 
we may have solutions of gases in gases, gases in liquids, 
liquids in liquids, solids in liquids, solids in solids, and so on. 
It is with one of the properties of solutions of solids or liquids 
in liquids (generally in water) that we shall deal in this 

li we examine .the structure of a plant by means of the 
microscope, we shall find that the plant-tissues consist of a 
number of small chambers, each of which with its contents 
is called a cell (Fig. 23). Each cell is bounded by a cell- wall 



of a substance called cellulose (a). Lining the wall is a layer 
of a semi-transparent gelatinous substance called protoplasm ; 
this layer is sometimes called the primordial utricle (6). The 
central portion of the cell consists of a watery liquid con- 
taining dissolved substances and called cell-sap (c). 

In 1854, PBINGSHEIM showed that if plant cells were placed 


Fio. 23. Single Plant Cell, seen FIG. 24. Diagrammatic Section of 
in Section. Plant Coll placed in Strong Salt 


in concentrated solutions of salt the protoplasmic lirnng shrank 
away from the cell-wall, the volume of the cell-sap becoming 
less. If these cells were now placed in pure water, the volume 
of the cell-sap increased again, and the shrunken protoplasmic 
lining returned to its original size and place. 

Thirty years later DB VBIES continued these experiments 
on a quantitative basis, taking as a standard the strength of 
a solution which was just sufficient to cause the protoplasmic 
lining to contract away from the cell-wall. By comparing 
the concentrations of solutions of different substances which 
were just able to produce this effect (" isotonic " solutions), 
he found that these concentrations were in the ratio of the 
chemical equivalents of the substances dissolved. 

These phenomena were apparently connected with those 
observed by DUTBOCHET (1726), the ABBE NOLLET (1748), 
PABBOT (1803), and FISCHEB (1822), on the passage of water 
through a parchment membrane into a solution of alcohol, 
cane-sugar or copper sulphate. If a solution of cane-sugar 
be placed in a thistle -funnel the mouth of which is closed by 
a piece of parchment or pig's bladder drawn tightly over it, 
and the funnel then immersed in a beaker of water, it will be 



FIQ. 25. 

found that the level of the liquid in the stem of the funnel 
gradually rises (Fig. 25). When a, solvent passes through a 
membrane as the water has clearly passed through in this 
case, the phenomenon is called osmosis, from the Greek 
a>(T,ao'g, a push. 

While water has passed through the membrane into the 
solution, no sugar can be de- 
tected in the water outside the 
funnel ; hence the membrane 
is semi-permeable that is, per- 
meable to the solvent but im- 
permeable to the solute. 

Up to the year 1867, the 
only semi-permeable membranes 
used were natural ones, such 
as pig's bladder in the above 
experiment and the protoplas- 
mic lining of plant cells in the 
experiments carried out by 
Pringsheim. But in 1867 TRATJBE showed that artificial semi- 
permeable membranes could be prepared, the best of which 
was found to be copper f errocyanide in the gelatinous state in 
which it is formed by precipitation from a solution of copper 
sulphate by addition of potassium f errocyanide solution. 

If a drop of a concentrated solution of copper sulphate is 
carefully introduced by means of a pipette into a dilute 
solution of potassium ferrocyanide, the drop becomes sur- 
rounded with a thin film of copper ferrocyanide. This acts 
as a semi -permeable membrane, and as the copper sulphate 
solution is concentrated and that of the potassium ferro- 
cyanide dilute, it will be found that the drop gradually 
increases in size by the passage of water through the mem- 
brane into the copper sulphate solution inside. 

PFEFFER made use of this artificial semi-permeable mem- 
brane to conduct quantitative experiments on osmosis. He 
published his results in 1877 in a work entitled Osmotische 
Untersuchungen. In Nollet's experiment with the thistle 
funnel, it will be remembered, a column of liquid was sup- 


ported, against the action of gravity, in the stem of the 
funnel. This column of water is therefore exerting a hydro- 
static pressure downwards, due to its weight. What is the 
force which balances this pressure, and keeps the column up ? 
The answer was given as " the osmotic pressure of the solution." 
This, of course, is not an explanation of the phenomenon, it 
is only a means of defining " osmotic pressure." The osmotic 
pressure in the above case will be equal in magnitude to the 
hydrostatic pressure of the column of liquid in the stem of 
the funnel. It was found that the osmotic pressures of even 
moderately concentrated solutions were very high, while 
those of very concentrated solutions often reached surprising 
values. Thus the osmotic pressure of a solution of cane-sugar 
containing 171 grams of sugar in 1 litre of water was found to 
be over 11 atmospheres. Now these great pressures rendered 
the use of the very delicate artificial membranes impossible, 
and that of the natural membranes difficult. Traube there- 
fore devised a plan for strengthening the copper ferrocyanide 
membrane. He soaked a porous pot in distilled water, then 
filled it with potassium ferrocyanide solution and placed it 
for some time in copper sulphate solution. The two solutions 
diffused into the wall of the pot and met in the middle, when 
a film of copper ferrocyanide was deposited in each pore. 
The pot was then washed in distilled water. The wall of the 
pot was now a sort of reinforced semi-permeable membrane 
of copper ferrocyanide, and the film deposited in this way 
was capable of withstanding high pressures. A further 
improvement was effected by MORSE in 1904, who used an 
electrolytic method for depositing the copper ferrocyanide in 
the pot. The membrane may also be strengthened by making 
it under pressure. 

If a porous pot prepared in this way is filled with the 
solution whose osmotic pressure is to be determined, closed 
with a stopper carrying a tube connected to a manometer or 
pressure-gauge, and then surrounded by water, the latter 
will pass in and the mercury in the manometer will rise 
(Fig. 26). After the rise has stopped the osmotic pressure 
may be read off the scale attached to the manometer. 



It is better in this experiment to add mercury to the open 
limb of the manometer until the pressure is just sufficient to 
prevent the water from entering 
the pot ; the osmotic pressure 
may then be read off as before. 
The great pressures developed, 
however, often burst this ap- 
paratus, and therefore other 
forms have been invented. The 
HARTLEY (1904) placed the solu- 
tion in a porous tube containing 
the copper ferrocyanide mem- 
brane in its walls, and enclosed 
the tube in a strong cylindrical 
vessel filled with water. Pressure 
was then applied to the solution 
by means of a narrow tube (Fig. 
27) until the level of the liquid 
in the latter showed no tendency 
either to rise or to fall, but 

remained steady. The pressure necessary to produce 
this equilibrium is equal to the osmotic pressure of the 


FIG. 26. Diagram of 

Apparatus for Measuring 

Osmotic Pressure. 


Water. S = Solution. 
M = Manometer. 

Porous Tube 

Strong Outer 

solution. This apparatus eliminates the danger of bursting 
the porous vessel. 



Accurate measurements of osmotic pressure enabled 
VAN'T HOFF in 1887 to show that the osmotic pressure of 
substances in solution- 

(i) was directly proportional to the concentration of the 
solution, for the same solvent and same solute, and 
at constant temperature ; 

(ii) was directly proportional to the absolute temperature ; and 

(iii) at constant temperature was the same for all solutions 

of non-electrolytes that contain the molecular weight 

in grams of the solute in the same volume of solvent. 

From (i) it follows that the osmotic pressure is inversely 

proportional to the volume of the solution. Hence we see 

that there is a very close analogy between gas-pressure and 

osmotic pressure, as the following table shows : 



P oc ~ if T is constant 

P oc T if V is constant. 

Equal numbers of molecules 
confined in equal volumes 
exert equal pressures if T 
is constant. [Follows from 
Avogadro's Hypothesis.] 

Osmotic P oc if T is constant. 

Osmotic P oc T if V is con- 

Equal numbers of molecules 
in the same volume of 
solvent exert equal osmotic 
pressures if T is constant. 

Further investigation proved that if the molecular weight in 
grams of a non-electrolyte is dissolved in 224 litres of solvent, 
the solution will exert an osmotic pressure of 760 mm. at C. 
Thus the analogy with gases is complete, for the G.M.V. of a 
gas at N.T.P. is 224 litres. 

We may therefore write for solutions, as for gases, 

PV = RT 

where P in this case is the osmotic pressure. This obviously 
gives us another method of determining the molecular weights 


of non-electrolytes, for the M.W. of such a substance will be 
the number of grams of it which when dissolved in 22-4 litres of 
a solvent will produce a solution having an osmotic pressure of 
760 mm. at <7. 

Example. 1-05 grams of a substance when dissolved in 112 c.c. 
of water gave an osmotic pressure of 940 mm. at 15 C. What is the 
Molecular Weight ? 

If 1-05 grams in 112 c.c. give a pressure of 940 mm. at 15 C. (== 288 

940 X 273 
Abs.), they will give a pressure of " mm. at C. 


.*. 1-05 grams in 22,400 c.c. will give a pressure of 
940 x 273 x 112 

288 x 22,400 

mrn. at C. 

. , . .940 x 273 X 112 * nor. u 

And since a pressure of - - -- mm. at C. is given by 

^oo X 

1-05 grams. 

.*. a pressure of 760 mm. is given by 

1-05 X 760 X 288 X 22,400 

940 X 273 X 112 
and this will be the molecular weight. 

In practice, however, the difficulties of determining osmotic 
pressures accurately are so great that this method of finding 
molecular weights is seldom used, since simpler methods are 
available. The cause of osmotic pressure and the reason 
for the semi-permeability of the membranes are unknown, 
though various theories have been advanced. 

It should be noted that, just as PV = RT is only approxi- 
mately true for gases, the deviations becoming considerable 
at high pressures, so is it only roughly true for osmotic 
phenomena. At high concentrations abnormalities become 
very marked. 
Phenomena related to Osmotic Pressure. 

1. Depression of the Freezing -Points of Solutions. Every 
pure liquid has a definite freezing-point. (It would be more 
correct to say melting-point, for on cooling a liquid it often 
happens that the temperature falls several degrees below 
the melting-point of the solid form before freezing begins. 
When freezing has once started, however, the temperature 


at once jumps up to the " true " freezing-point ; that is, 
the temperature at which the solid and liquid forms can exist 
in equilibrium. On the other hand, it is impossible to heat 
a solid above its melting-point without melting it. It should 
be understood, therefore, that in this connection, the " true " 
freezing-point, or melting-point, is meant.) If a substance is 
dissolved in a liquid, the freezing-point of the solution is 
lower than that of the pure solvent, and, for the same solvent 
and same solute, the extent of the lowering is directly pro- 
portional to the concentration. This is called Blagden's 
Law. Thus, if a 1 per cent, solution of a substance caused a 
depression or lowering of the freezing-point of 2-5, a 2 per 
cent, solution of the same substance in the same solvent would 
depress the freezing-point by 5-0, and so on. [The " depres- 
sion " is the difference, in degrees, between the two freezing- 
points.] There is obviously an analogy, then, between 
depression of the freezing-point and osmotic pressure, since 
both are directly proportional to the concentration. Further 
investigation confirmed this, and showed that the depression of 
the freezing-point of a solution was indeed directly proportional 
to its osmotic pressure. 

It will not surprise us, therefore, to learn that in 1883 
RAOULT showed that, for non-electrolytes, the lowering of the 
freezing-points of solutions of these substances in the same 
solvent was directly proportional to the molecular concentra- 
tions, since, as we have seen, the osmotic pressure is also 
proportional to the molecular concentrations. To make 
Raoult's discovery quite clear, let us consider the following 
example : 

Suppose that in each of five beakers, i, ii, iii, iv and v, we 
place 1,000 c.c. of water. The freezing-point of this water 
will be C. Let us now take the molecular weights in grams 
of the five following non-electrolytes 

Urea (= 60 grams) Alcohol (=46 grams) 

Cane-sugar (=342 grams) Resorcino} (=110 grams) 
Grape-sugar (=180 grams) 

and place the urea in i, the cane-sugar in ii, the grape-sugar 


in iii, the alcohol in iv, and the resorcinol in v, and stir 
until they have completely dissolved. If we now deter- 
mine the freezing-points of the five solutions, we find that 
in every case the freezing-point is 1-86C. Evidently 
we have here another method for determining the molecular 
weight of a non-electrolyte. All that we need do is to take 
another beaker with 1,000 c.c. of water in it, and find how 
many grams of the substance we have to dissolve in the water 
to get a solution that will freeze at 1-86 C. This number 
of grams will be the molecular weight of the substance. A 
consideration of Blagden's Law, however, will show us that 
since the depression in solutions of the same substance is 
directly proportional to the concentration, we need in practice 
only to dissolve a known weight of the substance in a measured 
volume of solvent, and note the depression caused. We can 
then work out the result by proportion. 

Thus, 0-346 grams of a substance dissolved in 10-3 c.c. of water 
gave a depression of the freezing-point of 0-21. Find the molecular 
weight of the substance. 

0-346 gms. dissolved in 10-3 c.c. produced depression 0-21 
/. 0-346 gms. ,, 1,000 ,, will give ,, 

, 0-21 x 10-3 /m . , _ . 
of j- (Blagdon's Law). 

0-21 x 10-3 

If a depression of - is caused by 0-346 gms. in 1,000 c.c., 


, i 0*0- , u ' 346 x 1<86 x 1 000 

then a depression of 1-86 is caused by -~~ r gms. 

U'Jl X lU-o 

= 298 gms. 
This is the molecular weight. 

The freezing-point or Cryoscopic method of determining 
Molecular Weights is extremely useful. For substances that 
are not soluble in water other solvents may be used, such as 
benzene, acetic acid, and, for metals, mercury. In books of 
tables are given the " Constants " for 100 grams of these and 
other solvents, tttat is, the depression of the freezing-point 
produced by dissolving the molecular weight in grams of any 
non-electrolyte in 100 grams of the solvent. The numbers 
for water, benzene, acetic acid, and mercury, are as follows : 



Benzene . 
Acetic acid 
Mercury . 

The depression for 

Depression* (K) for TOO 
grams of Solvent. 




100 c.c. can be worked out if the specific 
gravity of the solvent is known. 
Thus, the S.G. of benzene is 0-8784, 
therefore 100 grams of benzene occupy 
113-84 c.c., hence the constant for 100 
c.c. of benzene 

49 X 113-84 


= 55-8. 

FIG. 28. Freezing 
Point Apparatus. 

A. Tube containing solvent. 

B. Jacket tube. C. Thick 
glass jar. I). Beckmann 



It is necessary to say that Raoult's 
Law of Depressions holds only for 
dilute solutions, hence the above 
values are calculated from the results 
of dilute solutions, according to 
Blagden's Law. It can also be shown 
theoretically that K for 100 grams 

2T 2 
of a solvent = TTJTJJ- where T = the 

freezing-point of the solvent on the 
absolute scale, and L is its latent heat 
of fusion in calories. Thus for water 
T = 273 and L 79. 

We can now work out a general 
formula for this method of molecular - 
weight determination 
Suppose w grams of a substance in S 

gms. of solvent give a depression t 
.". 1 gram of a substance in S gms. 

of solvent gives a depression 

& L w 

/. M gms. (Molec. Wt.) of a sub- 
stance in S gms. of solvent 

a depression - X M. 
r w 


,* M grams of a substance in 1 gm. of solvent 

give a depression of X M X S 

M grams of a substance in 100 gms. of solvent give 

, . t . M . S __ 

a depression =- = K 

r w X 100 

' w X 100 

w . K . 100 

and M = 


Example. (a) Tammann showed in 1889 that 0-022 gms. of sodium 
dissolved in 100 gms. of mercury gave a depression of the freezing- 
point of mercury of 0-39. 

w.K.IOO ^ 0-022 x 425 X 100 
~~ *. S ~ 0-39 X 100 

= 23-8 

Therefore the molecular weight of sodium in mercury is 23' 8, i.e. 
here the molecular weight is the same as the atomic weight. Note 
that this gives us no information as to the molecular weight of solid 

(6) 0-53 gms. of a solid dissolved in 51-5 gms. of benzene lowered 
the freezing-point of the latter by 0-61. 

w . K . 100 0-53 X 49 X 100 

M = - 

t. 8 0-61 X 51-5 


N.B. As remarked before, it is much better to work out 
examples from first principles and not from a formula. Do 

not use the above formula, M = ' , until you know 

t . k5 

exactly what it means, how to get it, and how to work out 
your results without it. 

2. Lowering of the Vapour Pressure of Solutions. At a 
constant temperature, the vapour pressure of a liquid is 
constant. 1 If a substance is dissolved in the liquid, the vapour 
pressure of the solution is lower than that of the pure liquid, 
and it has been shown by TAMMANN and RAOULT that, for 
non -electrolytes as solutes, the G.M.W. of a substance dis- 
1 Provided that the vapour is in contact with its liquid. 


solved in a constant weight of the same solvent produces a 
constant relative depression of the vapour ^pressure. The 

f f 
" relative " depression is - ~- y where /is the vapour pressure 

of the pure solvent and /' that of the solution. Raoult found 

by experiment that the formula that agreed best with his 

f & n 

observed results was f = -^ . > where n = number of 

gram-molecules of solute and N = number of gram-molecules 
of solvent. Hence by vapour-pressure measurements we can 
calculate molecular weights. 

Example. 3- 19 gins, of a substance dissolved in 115 gms. of benzene 
gave a solution of vapour pressure 731*5 mm. The vapour pressure 
of pure benzene at the same temperature was found to be 750 mm., 
and the molecular weight of benzene = 78. Find the molecular weight 
of the substance. 

/ _ /' 750 - 731-5 18-5 






** N -f- n 




/ IN -f- n 

n 18-5 

1-47 + n 750 

/. 750 n = 27-2 + 18-5 n 

/. 731-5 n = 27-2 
/. n = 0-037 


But n = 

M. Wt. 

Molecr. Wt. of Substance 

_ 3 ' 19 
~~ (H)37 

= 86. 

Vapour pressures are difficult to determine accurateJy in 
practice, hence this method is seldom employed, but since 
lowering of vapour pressure is, over short ranges, roughly 
proportional to the elevation of the boiling-point thereby 



produced, an ebullioscopic method of finding molecular weights 
has been elaborated. 

Liquids boil when their vapour pressures become equal 
to the external pressure, the normal value of which is 760 
mm. In Fig. 29 AB represents the vapour pressure curve 
of the pure solvent, CD and EF the V.P. curves of two solu- 
tions, EF being that of the more concentrated. 

When the vapour pressures of the liquids reach 760 mm. 
they will boil, but the solvent will reach this pressure at a 


-760 mm. 

FIG. 29. 

lower temperature (t) than the first solution (f) and this 
again at a lower temperature than the more concentrated (t"), 
t t t', and t" represent the boiling-points of the solvent, first 
solution, and second solution respectively. Now it is found 
experimentally that over short ranges of temperature, the 
curves AB, CD, and EF are practically parallel. The vapour 
pressures at t are tB for the solvent, tNL for the first solution, 
and JE for the second. Therefore the relative lowering of the 
vapour pressures are, for the first solution, 


BM\ , *B-E/ BE 

= IB > and for the second 

Therefore the ratio of the relative lowerings for the two 

. BM / BE BM _ A . A _ ^ , 
solutions is -^- / -75- = T5. But since AB, CD, and LP 

4.' 11 11 1 BM ' .'11 1 4. BD T 

are practically parallel, ^ is practically equal to -^-^. In 

other words, since BD and BF are equal to t' i and i" t 
respectively, the relative depressions of the vapour pressures 
of these two solutions are in the same ratio as the elevations 
of their boiling-points (t' t and t" t). 

We can, therefore, use the elevation of the boiling-point 
of a solution exactly as we use the depression of the freezing- 
point, for purposes of molecular weight determination. The 

i i j *. xi ^ w.K.100 

formula, worked out in exactly the same way, is M= - ~ , 


where w is the weight of substance taken, K the constant 
elevation of the boiling-point for 100 grams of the solvent, 
t the observed elevation, and S the weight of solvent used. 
The constants for 100 grams of different solvents are here 

K, Constant for 

Solvent. 100 Grams. 

Water ........ 5-2 

Alcohol ........ 11-6 

Ethor ........ 21-0 

Benzene ........ 25-7 

2T 2 

Here also K = T , where T is the boiling-point of the 

pure solvent in degrees absolute, and L is the latent heat of 
vaporization of the solvent, in calories. 

An apparatus used for the determination of molecular 
weights by boiling-point elevation is shown in Fig. 30. 

General Remarks. The osmotic pressure laws, and the 
relationships between the molecular concentrations of solu- 
tions and the depressions of the freezing-point and vapoar 
pressure and elevation of the boiling-point, hold good only 



for non-olectrolytes and dilute solutions. We shall see later 
on (Chap. XIII) that the exceptions and irregularities that 
were discovered in the case of electrolytes gave rise to a great 


Fio. 30. Boi ling-Point Apparatus. 

F. Flask for boiling liquid. N. Graduated Boiling Tube. H, C. Exit for vapours. 
T. Thermometer. 

deal of theoretical speculation, much of which has been 
supported by experimental results. 


1. What do you understand by the osmotic pressure of a solution ? 

2. Describe a method of measuring osmotic pressure. 


3. Compare and contrast osmotic pressure with gas pressure. 

4. What is the principle of Raoult's cryoscopic .method of deter- 
mining molecular weights ? 

6. What physical properties of a solution are related to the osmotic 
pressure of the solution ? Of what practical value is the relation- 
ship ? 

6. 0-6418 gm. cane sugar is dissolved in water and made up to 56-8 
c.c. at a temperature of 15 C. Find, to the nearest mm. of mercury, 
the osmotic pressure produced. [Mol. wt. of cane sugar is 342.] 

7. Find the osmotic pressure, in atmospheres, produced by a 13 
per cent, solution of ammonia at 16 C. [Mol. wt. ammonia = 17.] 

8. Find the percentage strength of a solution of cane sugar which 
at 7 C. is isotonic with a 2 per cent, solution of dextrose at 27 C. 
[Mol. wts. Cane sugar, 342. Dextrose, 180.] 

9. 0-4277 gm. of an organic compound is dissolved in water and made 
up to 100 c.c. at 12 C. The osmotic pressure produced is equal to 
380 mm. of mercury. Find the molecular weight of the compound 
to the nearest whole number. 

10. At what temperature is a 2 per cent, solution of glycerol isotonic 
with a 4-5 per cent, solution of glucose at 15 C., and what is the osmotic 
pressure in atmospheres of the isotonic solutions ? [Molecular wts. 
Glycerol, 92. Glucose 180.] 

11. Find, to the nearest atmosphere, the osmotic pressure of an 
aqueous solution of an organic compound, at its freezing-point, which 
is 0-33 C. The molecular depression of the freezing-point of water 
is 18-6. 

12. 0-3132 gm. fructose was dissolved in 14-5 gms. water. The freez- 
ing-point of the solution was found to be 0-223 C. The molecular 
depression for water is 18-6. Find the molecular weight of fructose 
to the nearest whole number. 

13. The freezing-point of a solution of 0-4813 gm. acetone in 42-8 gms. 
acetic acid is found to be 15-84 C. The freezing-point of pure acetic 
acid is 16-60 C., and its molecular depression 39. Find the molecular 
weight of acetone. 

14. A benzene solution containing 2 gms. of an organic compound to 
100 gms. of the solvent is found to freeze at 4-07 C. Given that the 
molecular weight of the compound is 120 and the freezing-point of 
pure benzene 4-90 C., find the molecular depression of benzene. 

15. A solution of 3 gms. of an organic compound in 100 gms. of phenol 
freezes at 38-94 C. Given that the molecular weight of the compound 
is 220 and the molecular depression of phenol 73-8, find the freezing- 
point of pure phenol. 

16. It is found that an aqueous solution of hydrogen peroxide 
freezes at 0-35 C. Find the weight of hydrogen peroxide present 
per 100 gms. of water, given that the molecular depression of water is 
18-6. [Molecular weight of hydrogen peroxide = 34.] 


17. Given that the fall in freezing-point of a solution containing 
0-9 152 gm. of an organic compound of molecular weight 59, in 25-0 gms. 
benzene is 2-42, find the molecular weight of another compound, 
I 5480 gms. of which when dissolved in 48-0 gms. benzene lower the 
freezing-point by 0-91. 


Chemical Affinity. It is characteristic of human nature 
to ask why a thing happens before investigating how it 
happens. Yet it is by finding out the exact way in which 
a thing happens that we are most often enabled to find out 
the reason for its occurrence. Chemists, however, being men 
of like passions with ourselves, began to make theories about 
the cause of chemical change long before they were in a 
position even to investigate thoroughly the progress of 
chemical change. EMPEDOCLES (ca. 500 B.C.) assumed that 
chemical changes were brought about by the love or hate 
which various substances had for one another ; two sub- 
stances that " loved " one another would unite to form a 
third substance, while decompositions represented the 
separation of substances that hated one another. As a 
matter of fact, we know very little more about the cause of 
chemical reaction at the present day than Empedocles did, 
but we conceal our ignorance in a much more imposing way. 

In the Middle Ages, it was considered that substances 
which would combine together were more or less similar or 
related to one another mercury, for example, readily dis- 
solves metals and in order to express this relationship the 
word affinity (from the Latin affmitas, relationship) was used. 
Substances therefore combine with one another on account 
of their chemical affinity. It was observed, however, by later 
chemists that in general it is substances which are most opposed 
in qualities that most readily react together, e.g., sulphur and 
iron, but the term affinity by then had become well estab- 



lished, and has remained to the present day as the name for 
that chemical force which brings about reaction. The exact 
nature of this force still remains a mystery. NEWTON sup- 
posed that it might be gravitational ; others have considered it 
to be electrical a view which is now generally accepted. The 
electrical theory was warmly supported by BERZELIUS at the 
beginning of the nineteenth century, who pointed out that 
in a large number of cases it was elements of dissimilar electro- 
chemical character which most readily combined together. 

The " affinity " of acids for alkalis had of course been 
observed very early, and MAYOW (1674) explained the libera- 
tion of ammonia from sal-ammoniac, by the action of potash, 
by the greater affinity that the acid possessed for potash 
than for ammonia. Further elaboration of the theory led to 
the production, by GEOFFROY (1718) and BERGMANN (1775) 
of Tables of Affinity, in which substances were arranged in 
their supposed " order of affinity, " or the order in which they 
would turn one another out of combination in a particular 
reaction. Thus copper will be precipitated if iron is placed 
in a solution of copper sulphate, therefore the affinity of iron 
is greater than that of copper, in this reaction. In general, 
it was supposed that if the element A has a greater affinity 
for C than B has, then A will decompose the compound BC 
completely, forming AC with liberation of B. 

In the light of this theory, let us consider the reaction be- 
tween steam and iron. If steam be passed over heated 
iron, iron oxide and hydrogen are formed, therefore the 
affinity of iron for oxygen must be greater than that of 
hydrogen for oxygen, for the iron has turned the hydrogen 

3Fe + 4H 2 O = Fe 3 4 + 4H a . 

oxide of iron. 

But if we pass hydrogen over heated iron oxide we shall 
find that steam* is formed and iron left. Therefore the 
affinity of hydrogen for oxygen is greater than that of iron, 
which is the opposite conclusion to that at which we arrived 
before ! 


Similar phenomena were noticed by BERTHOLLET in 1801, 
who deduced from his results that there must be another 
important factor in chemical reaction, and suggested that 
this factor was the concentrations of the substances taking 
part. Thus in the above experiment, in the first case the 
concentration of the steam is great and that of the hydrogen 
negligible, since the latter is carried off as soon as formed, by 
the current of steam. In the second case, the hydrogen is in 
great concentration and the steam formed gets swept away. 
Berthollet's Law, then, may be stated in these words : 
In a chemical reaction, each substance will react according to 
(a) its affinity, (b) its concentration. 

Chemical Equilibrium. We have seen that the paradox 
of the reactions between iron, iron oxide, steam and hydrogen 
can be explained by the fact that in one case the hydrogen 
and in the other the steam are swept away from the sphere 
of action. What will happen if some iron and steam are 
heated together in a closed vessel ? Experiments made to 
test this showed that at first iron oxide and hydrogen were 
formed, but that after a time no further change could be 
detected, although some of the iron and some of the steam 
were still left. However long the vessel was kept, at the 
same temperature, the composition of the mixture inside 
remained unaltered ; in other words, a state of equilibrium 
has been reached. As, however, steam and iron, and iron 
oxide and hydrogen, are present, chemical reactions must 
still be going .on, and the equilibrium is therefore explained 
by assuming that as much steam is formed in a given time as 
is decomposed in that time ; that is, the speed of the one 
reaction is equal to the speed of the other, and the equilibrium 
is a kinetic equilibrium. 

The speed of a chemical change may thus be defined as 

weight of substance changed 

equal to -. --; . 

^ time taken 

Reactions which, like that represented by the equation 

3Fe + 4H a O = Fe 3 4 + 4H 2 , 
can proceed either way, are called reversible reaction*, and it 


is usual in such cases to replace the sign of equality by the 
" reversed arrows " sign, s s , e.g., 

3Fe + 4H 2 ^z Fe 3 4 + 4H 2 . 

When the conditions are so arranged that both the forward 
and the backward reactions are proceeding at the same speed, 
and equilibrium is thus established, the action is called a 
balanced action. 

Many reactions are reversible, and it is possible that if we 
could produce the right conditions, all would be so ; but 
since in the present state of our knowledge, " all the king's 
horses and all the king's men could never put " the products 
of the explosion of trinitrotoluene " together again," it is 
convenient to distinguish between reversible and irreversible 
Examples of reversible reactions. 

(i) If concentrated hydrochloric acid is poured on anti- 
mony sulphide, sulphuretted hydrogen and antimony chloride 
are formed 

Sb 2 S 3 + 6HC1 = 2SbCl 3 + 3H 2 S. 

If, however, sulphuretted hydrogen is passed through a 
solution of antimony chloride, the reverse reaction takes 
place, and antimony sulphide is precipitated. 

(ii) If a mixture of 2 volumes of hydrogen and 1 volume 
oxygen is ignited, an explosion occurs and steam is formed 
2H 2 + 2 - 2H 2 0. 

But if a stream of steam is passed over a white-hot platinum 
wire, some of the steam is split up again into hydrogen and 

(iii) " Dissociation " is a reversible reaction (Chap. IX). 

(iv) The action of an organic acid upon an alcohol to form 
an ester, e.g., acetic acid upon ethyl alcohol 

CII 3 .COOH + C 2 H 6 OH ^= CH 3 .COOC 2 H 5 + H 2 O. 

Ethyl acetate. 

If ethyl acetate is heated with water, the reverse change 
slowly occurs until equilibrium is set up. 

(v) The decomposition of calcium carbonate by heat 
CaC0 8 ^= CaO + C0 a . 


Guldberg and Waage's Law or The Law of Mass Action. 

In 1850 WILHELMY showed that the rate of a chemical 
change at a given moment is directly proportional to the 
concentration of the reacting substances (Wilhelmy's Law). 
The work of Berthollet and Wilhelmy was generalized and 
put into a more definite form by GULDBEKG AND WAAGE, as 
the Law of Mass Action. 

This states that the rate of a chemical change at a given 
moment is directly proportional to the concentrations at that 
moment of the substances taking part in the change. The 
" concentration " for this purpose is usually expressed in 
gram-molecules per litre and is then called the active mass 
of the substance. 

In its general form, then, the Law of Mass Action may be 
expressed as follows. If n molecules of A, of active mass M, 
n' molecules of B, of active mass M', n" molecules of C, of, 
active masp M", etc., take part in a chemical change, then 
the rate of the reaction at the moment when the active masses 
have the above values is proportional to 

M n X M' n ' x M" n " X etc. 

Consider the action M + N = P. 

Let the active mass of M be a and that of N be 6. Then 
the rate of change of M and N into P will be proportional to 
a. b y or equal to a.b x a constant, that is, 
rate of above reaction = k.ab. 

Similarly the rate of the reaction 2M + N = R will be 
k.a 2 b at the moment when the active mass of M is a and 
that of N is 6. 

If we have a reversible reaction, such as A + B * C + D, 
it is clear that equilibrium will be set up when the rates of 
the two opposing reactions are equal. Let the active masses 
of A, B, C, and D at equilibrium be respectively a, 6, c and d. 
Then the rate of the forward reaction will be k.ab and that 
of the reverse reaction k' .cd. These, however, are equal. 

Hence k.ab = k' .cd, or -5 = K. 

It follows that if we now increase the concentration of 


A or B the equilibrium should be shifted in such a way that 
more C and D are formed, a deduction which is very neatly 
proved by GLADSTONE'S experiment. When ammonium 
thiocyanate solution is added to ferric chloride solution, a 
blood-red colour is produced owing (it is said) to the forma- 
tion of ferric thiocyanate 

FeCl 3 + 3NH 4 CNS ^= Fe(CNS) 3 + 3NH 4 C1. 

This action is reversible, as shown in the equation. Now 
if the Law of Mass Action and the deduction we have made 
from it are true, when we add ferric chloride to the liquid, the 
result should be an increase in the amount of ferric thiocyanate 
present, and therefore the liquid should become of a deeper 
tint. This is found to be the case ; addition of ammonium 
thiocyanate produces the same result ; addition of ammonium 
chloride makes the colour paler y because, the active masses 
on the right-hand side having been thus increased, the 
equilibrium is shifted in the direction of ferric chloride and 
ammonium thiocyanate. 

To return to the action of iron on steam. Here we have 
an action, 

3Fe + 4H 2 ^= Fe 3 4 + 4H 2 , 

that takes place between two substances which are in 
different physical states, one a solid and the other a gas. If 
at any moment the active mass of the iron is / and that of the 
steam s, then the velocity of the forward reaction will be 
&./ 3 s 4 , where & is a constant. But the iron can react with 
only that part of the steam which reaches its surface, while, 
owing to the movement of its molecules, any part of the 
steam is potentially capable of getting into contact with the 
iron. Now if we have a large surface of iron exposed, it is 
obvious th&t more steam may be in contact with it at a given 
moment ; but at the same time, this large surface of iron will 
also give a proportionately larger surface of iron oxide to be 
decomposed by tne hydrogen present. What is the con- 
clusion to be drawn from this ? It is surely this, that since 
any variation in the amount of solid present in a gaseous 
reversible reaction will affect the rates of both reactions equally, 


the active mass of the solid may be taken as constant, whatever 
amount of it is present. 

Thus, in the reaction 3Fe + 4H 2 - ^ Fe 3 4 + 4H 2 , let 
/ be the active mass of the iron, s that of the steam, c that of 
the iron oxide, and d that of the hydrogen, at equilibrium. 

Then k.f*.s* = F.c.d 4 . 

f and c, however, represent the active masses of solids and 
are therefore constant. 

4 8 

.". "71 is constant, whence -7 is constant = K. 
d* d 

This means that the percentage composition by volume of 
the gaseous mixture of steam and hydrogen which is in 
equilibrium with iron and iron oxide at a constant tempera- 
ture is itself constant. Experiment shows that this is so. 

K is called the equilibrium constant of the reaction. 

Ester -formation. Suppose we mix 1 gram-molecule of 
acetic acid with 1 gram-molecule of alcohol. A reaction 
occurs according to the equation 

CH 3 .COOH + C 2 H 6 OH ^= CH 3 .COOC 2 H 5 + H 2 0. 

Suppose that at a time t equilibrium has been set up, and 
that a grarn-molecules of the acid have been transformed into 

ester. Let the volume of the liquid be V. Then the active 

j ^ 

mass of the acid at equilibrium = - v , and it follows from 

the equation and from the fact that we started with 1 gram- 
molecule each of alcohol and acid, that the active mass of 

the alcohol is also ^ . Similarly, the active masses of the 

ester and water will be ^ and ^. But we have assumed 
that equilibrium has been set up 

, 1 a 1 a __ 7 , & a 

" * ~v~ ' ~~V~ = * V ' V 

. *' _ a - a) 2 nr (i - ) 8 __ 

" ~~~~ or ~ -** 


Hence by finding the weight of acid left at equilibrium we 
can calculate K, 

By mixing acetic acid and alcohol in other proportions we 
can arrive at a general formula. In this example, only one 
molecule of each substance is concerned, so if we let C l9 C a , 
C 3 , and C 4 represent the active masses of acid, alcohol, ester, 
and water respectively at equilibrium, we get 

A;.C 1 .C 2 = ^.C 3 .C 4 
,,. k' active mass of acid X active mass of alcohol 

QJ IT ___ _ , - . _. _ ___ 

k active mass of ester X active mass of water 

Analysis of various equilibrium mixtures of these four sub- 
stances will enable us to test the theory by observing whether 
the values for K actually obtained are really constant. 
Experiments have shown that this is so. The value of K 
here is about 0-25. 

Dissociation of Hydrogen Iodide. It will be remem- 
bered that hydrogen iodide dissociates on heating. 

2HI ;== H 2 + I 2 , 

and that since no change in volume occurs we concluded that, 
according to Le Chatelier's principle, pressure changes should 
have no effect on the degree of dissociation at a constant 
temperature. This can be shown also by the Law of Mass 

Suppose that at a certain temperature equilibrium has 
been attained in the above action. Let the active mass of 

hydrogen iodide be ^, and that of hydrogen ^, whence that 
of iodine must clearly also be ^. Then 

the V's cancelling out. Hence the proportion of hydrogen 
iodide dissociated is independent of the volume, and therefore 
of the pressure, if the temperature is constant. In cases 
where the V's do not cancel, we should expect to find that 


changes of pressure would affect the degree of dissociation 

Nitrogen Peroxide. N 2 4 ~ ^ 2N0 2 . Suppose thai 
equilibrium has been set up, and that the volume of the gas 
is V. Let there be a gm. molecules of N 2 4 and 6 gm 
molecules of N0 2 . Then the active mass of N 2 4 i* 

^, and that of the N0 2 ^. 

(since there are TWO N0 2 molecules concerned). 

*L-?L Z. Z 
* & ~~ V X 6 X 6 

Here, then, the degree of dissociation will depend upon the 
volume and therefore on the pressure. 

Phosphorus Pentachloride. PC1 5 ; * PC1 3 + C1 2 . At 
equilibrium let the volume of the gas be V, the number oJ 
gram-molecules of phosphorus pentachloride a, and thai 
of phosphorus trichloride and chlorine b of each. Then the 

respective active masses are ^, ^, and ^, and, as in the 

case of nitrogen peroxide, r- = ,' 2 , or the degree of dis- 

sociation will depend upon the pressure. Now Wurtz showed 
that if phosphorus pentachloride is heated in an atmosphere 
of chlorine, the dissociation of the former was reduced. Does 
the law of mass action enable us to understand this ? Take 
the above equation 

k a ~-k f b - 

* y ~~ * v * y 

Suppose we add more chlorine, keeping the volume constant. 

a _ 6 (6 + c) 



where c = number of additional gram-molecules of chlorine 
introduced. It is clear that b(b + c) is greater than 6.6, 
therefore to maintain equilibrium a must have increased, 
that is, more undissociated phosphorus pentachloride will 
have been formed and the degree of dissociation thereby 

But suppose we add more chlorine at the same con- 
centration, say nV volumes. Then 

If a = V b b ( n + *) 
V + nV V + nV * V + nV 

. , a 7/ 6 6 

whence fc . = AT , ^ . ^, 

which is the same equation as before the addition of the 
chlorine ; hence, adding chlorine at the same partial pressure 
has no effect on the degree of dissociation. To obtain his 
results, therefore, Wurtz must have added the chlorine with- 
out allowing the volume to increase. 


1. Write an account of chemical affinity. 

2. What do you understand by a reversible reaction T Give examples. 

3. State the Law of Mass Action. Discuss the equilibrium between 
iron, iron oxide, steam, and hydrogen in terms of this law. 

4. What would be the composition of the equilibrium mixture 
finally formed if 1 gram-molecule of ethyl alcohol were added to 4 
gram-molecules of acetic acid ? K = 0-25. 

5. Discuss the equilibria 

(i) 2HI ^ H 2 + I 2 . 
(ii) PC1 5 ^PC1 ( , -f C1 2 . 
(lii) CaO + CO 2 ^ CaCOg. 


It will be remembered that the laws of osmotic pressure 
and the related effects (such as depression of the freezing- 
point, elevation of the boiling-point, etc.) were found to be 
true as far as non-electrolytes were concerned. When, how- 
over, electrolytes are used, discrepancies and anomalies 
occur, of such a nature that they at once recall the " abnor- 
mal " vapour densities of ammonium chloride and other 
substances. Thus, if we determine the molecular weight of 
common salt by the depression of the freezing-point, we shall 
find that in dilute solution the value obtained is not 58-5, 
as would be expected from the formula NaCl, but 29-25. 
Potassium nitrate treated in the same way gives a molecular 
weight of 50-5, whereas KN0 3 = (39 + 14 + 48) = 101 ; 
sulphuric acid, again, by the cryoscopic method gives a 
molecular weight not of 98 but of approximately 33. If 
more concentrated solutions of these substances are taken, 
the molecular weights determined by the cryoscopic or osmotic 
pressure methods gradually rise, finally in very concentrated 
solutions approaching more or less closely the values calcu- 
lated from the formulae. 

We are at once reminded of the phenomena shown by 
nitrogen peroxide. Here at a temperature of 22 the vapour 
density is 46 and the molecular weight therefore 92, corre- 
sponding to the formula N 2 4 . As the temperature rises, the 
vapour density falls till a constant value of 23 is reached, 
corresponding to the formula N0 2 . This phenomenon we 
explained by assuming that dissociation had taken place ; 



can we explain the abnormal results obtained above in this 
way ? Apparently not, for it is difficult to see how sodium 
chloride could dissociate ! VAN'T HOFF therefore proposed 
to shelve the difficulty for the time being by frank admission 
that electrolytes were exceptions to the osmotic-pressure 
laws, and introduced the use of a coefficient, i, to represent 
the ratio of the observed value of the depression of the freezing- 
point, elevation of the boiling-point, osmotic pressure, or 
lowering of the vapour pressure, to the corresponding value 
calculated from the generally-accepted formula of the electro- 
lyte used. Thus, we should expect 58-5 gms. (the G.M.W.) 
of sodium chloride, when dissolved in 10,000 c.c. of water, 
to give a solution freezing at 0-186 C. The actually 
observed freezing-point of the solution is, however, 0-37 C. 
In this case 

i is therefore the number by which one has to multiply the 
value for the osmotic pressure, etc., calculated from the 
formula, in order to get the observed value, or 
P observed 
P calculated 

The osmotic pressure equation will now become 
PV = i KT 

Another point that had puzzled chemists for some time 
was that when the gram-equivalents of many inorganic, and 
some organic, acids were neutralized by addition of the 
gram -equivalent of either potash or soda, the number of 
calories evolved was roughly the same in every case. This 
was expressed by saying that the heat of neutralization of 
acids is approximately constant (about 13,800 calories). The 
actual figures for a few acids are given below 

H eat of Heat of 

Acid. Neutralization. Acid. Neutralization. 

Nitric . . . 13,700 Hydriodic . . 13,700 

Hydrochloric . 13,700 Oxalic . . . 14,100 

Sulphuric . . 15,600 Citric . . . 13,100 
Hydrobromic . 13,800 


This seemed a very significant fact, but nq satisfactory 
explanation of it was forthcoming. 

Meanwhile, the study of electrolysis had not been neglected. 
We may recapitulate here the main features of electrolysis as 
formulated by Faraday. 

(i) The substances that appear at the electrodes are called 

(ii) The passage of the same quantity of electricity through 
solutions of electrolytes will liberate weights of different ions 
in the ratio of the chemical equivalents of those ions ; in other 
words, the electrochemical equivalent of a substance is equal 
to its chemical equivalent. 

(iii) The gram -equivalent of any ion carries 96,000 coulombs. 
From this it follows that each atom or group of atoms form- 
ing the particles of an ion must carry a number of unit 
charges equal in number to its valency, if the charge carried 
by a hydrogen atom is taken as unit. For 1 gram of hydro- 
gen must contain the same number of atoms as 63*5 grams 
of copper, since 63-5 is the atomic weight of copper. But 
the electrochemical and chemical equivalent of copper is 
31-8, therefore 31-8 grams of copper and 1 gram of hydrogen 

carry 96,000 coulombs each, or -^ atoms of copper carry as 

much electricity as n atoms of hydrogen, .'. 1 atom of hydro- 
gen carries half as much electricity as 1 atom of copper. 
If therefore we take the charge carried by a hydrogen atom 
as unit, the copper atom will carry 2 charges ; now the 
valency of copper is 2. Similarly, the valency of the S0 4 
group is 2, therefore the S0 4 group carries 2 unit charges. 

Another well-known fact was that the products of electro- 
lysis appear simultaneously at the two electrodes, however 
far apart these may be. Hence, in the electrolysis of dilute 
sulphuric acid, for example, a hydrogen atom liberated at a 
particular moment at the cathode may not belong to the 
same molecule of sulphuric acid as the S0 4 group which is 
discharged at the anode at the same moment. 

To explain these phenomena, various theories had been 
suggested, but they all assumed that the actual disruption 


of the molecules of the solute was caused by the current. If 
this were so, then a current would have to do work during 
electrolysis, to overcome the chemical forces holding the 
parts of the solute molecules together ; now experiment 
showed that no more work is expended by the current in 
passing through a solution of an electrolyte than it would 
expend in passing through a wire of the same resistance, when, 
of course, no electrolysis takes place. 

All these apparently unrelated phenomena of " abnormal " 
osmotic pressure, heats of neutralization, and electrolysis, 
were shown to be explicable by a general hypothesis, due to 
the Scandinavian chemist ABRHENIUS and called the Ionic 
Theory or Theory of Electrolytic Dissociation ( 1 887 ) . Arrheniua 
supposed that when an electrolyte is dissolved in water it 
splits up completely in very dilute solution and to a less 
extent in more concentrated solution, into charged atoms or 
groups of atoms which, by a transference of Faraday's word, 
were called ions. Each ion is assumed to exert an osmotic 
pressure equal to that exerted by an un-ionized molecule ; 
that is, as far as osmotic pressure and the related phenomena 
are concerned, ions and molecules produce the same effect. 

This theory met with a great deal of opposition at first, 
partly owing to a misunderstanding. Critics objected that 
it seemed ridiculous to suppose that free sodium atoms and 
free chlorine atoms were present in a solution of common 
salt ! Arrhenius, however, never made such a supposition. 
He suggested that not ordinary sodium and chlorine atoms 
were present in the free state in a solution of common salt, 
but very highly charged sodium and chlorine atoms. Now 
there is no difficulty in assuming that these highly charged 
atoms, or ions, might be very different in behaviour from the 
same atoms uncharged, and when this point was made clear 
opposition became less. Even now, however, the ionic theory 
is not wholly without objection, though with the advance of 
our knowledge of the electrical structure of matter it Jbas 
immeasurably strengthened its position and undoubtedly 
represents a triumph of man's insight into the phenomena 
of Nature. 


The most serious difficulty which the ionic theory has to 
face, and which up to the present has received no entirely 
satisfactory explanation, is the source of the energy required 
for ionization. When 23 grams of sodium combine with 
35-5 grams of chlorine to form 58-5 grams of sodium chloride, 
97,800 calories are evolved. To ionize 58-5 grams of sodium 
chloride it would seem that in addition to the energy repre- 
sented by the heat of formation one must supply the extra 
energy required to overcome the electrostatic attraction of 
the charged sodium and chlorine ions for one another. 

A solution of an electrolyte is electrically neutral because 
it contains equal numbers of positive and negative ions. To 
represent an ion carrying one positive charge, a dot is placed 
after the symbol for the uncharged atom or group of the ion, 
thus Na* represents an ion of sodium and NH 4 * an ion of 
ammonium. Fe" and Fe"" represent a ferrous ion with 2 
positive charges and a ferric ion with 3 positive charges respec- 
tively. In the same way dashes are used to indicate negative 
charges ; thus Cl' is a chlorine ion carrying one negative 
charge and S0 4 " a " sulphate " group carrying 2 negative 

In a fairly concentrated solution of an electrolyte, some of 
the solute molecules are ionized and others remain un-ionized, 
and the fraction of the whole that exists as ions is called the 
degree of ionization of the electrolyte at the particular con- 
centration (or " dilution ") used. We may imagine that in 
the solution molecules of the electrolyte are continually being 
ionized and that ions are continually recombining to form 
un-ionized molecules. In other words, there is a state of 
kinetic equilibrium between ions and un-ionized molecules, 
e.g., in a solution of salt 

NaCl ^ Na' + Cl'. 

The degree of ionization will be increased by further dilution 
and becomes practically equal to 1, with most electrolytes, 
at a dilution of about a thousandth normal, N/1,000. A 
degree of ionization of 1 means that the whole of the solute 
is ionized ; this clearly could theoretically occur only at 


infinite dilution. However, the degree of ionization at a 
dilution of N/1,000 so closely approaches to 1 that the differ- 
ence is usually negligible. 

Calculation of Degree of Ionization. Two independent 
methods of determining the degree of ionization of a solute 
in solution will be described. As they give more or less 
concordant results, the probability of the " truth " of the 
ionic theory is thus rendered greater. 

(i) Suppose that we take 1 gram-molecule of an electro- 
lyte which would give 2 ions per molecule, i.e., a " binary " 
electrolyte, such as salt, potassium nitrate, or hydrochloric 

Let the degree of ionization be x. 

Then we should have in solution 1 x undissociated 
molecules and 2x ions. But molecules and ions behave in the 
same way as far as osmotic pressure and related effects are 
concerned ; therefore the ratio of the observed osmotic 
pressure (etc.) of such a solution to the osmotic pressure (etc.) 
calculated on the assumption that 1 gram-molecule of the 
electrolyte has been dissolved without ionization would be 

J _ X -4- 2x 

- . This ratio, however, is van't Hoff's coefficient i 

(p. 127). We may therefore write 

or x = i 1. 

Thus in the case of a binary electrolyte we may find the 
degree of ionization as follows : 

(i) Calculate the theoretical depression of the freezing. 

point for the solution (d), from the formula of the 

(ii) Find the actual depression (d') of the freezing-point for 

the Solution, experimentally. 
Then the degree of dissociation of the electrolyte at this 

concentration will be i 1 or -3- 1 



For an electrolyte that forms n ions per molecule the 
expression will become 

1 x + nx 

$ =r= 



(iii) The conducting power or " conductivity " of a solution 
will be directly proportional to (a) the number of ions present 
per c.c. and to (6) the speed with which they move. For 
different solutions of the same solute in the same solvent we 
may legitimately assume that (6) is constant at constant 
temperature. Hence the conductivity will be proportional 
to the number of ions per c.c. This will be increased by 
dilution in so far as the degree of ionization is raised, but will 
be decreased by dilution in so far as the solution becomes 
weaker. We can allow for the second of these factors in the 
following way. The specific conductivity of a solution is the 
conductivity between two parallel faces of a centimetre cube 
of it. Suppose that for a solution containing 1 gram-molecule 
of an electrolyte in g litres the specific conductivity is 8. If 
we dilute the solution with an equal volume of water, and 
suppose for the moment that the degree of ionization remains 
the same, there will be only half as many ions per c.c. in the 
diluted solution as in the origina solution. Hence the 

specific conductivity will now be - . But if we take the 

molecular conductivity as the specific conductivity of the 
solution multiplied by the number of litres of the solution 
containing 1 gram-molecule of the electrolyte, then the 
molecular conductivity of the first solution will be s X g, say 

m, while that of the second will be - x 2g, or s x gr, i.e., 

again m. In other words, by taking the molecular conduc- 
tivity instead of the specific conductivity, we cut out any 
change due merely to dilution. Any change in molecular 
conductivity on dilution must therefore be due to a change 
in the degree of ionization. 


[f }, = molecular conductivity of a solution and a its 
degree of ionization, 

then A varies as a or A = a X a constant. 

The value of the constant can be calculated from the 
conductivity at " infinite " dilution (practically, at a dilution 
of N/1,000), for here a = 1 and A then = the constant. 

Hence, if x = the degree of ionization of a solute in a 
solvent at a given dilution, h d = the molecular conductivity 
of this solution, and ^ = the molecular conductivity of an 
"infinitely " dilute solution of the solute in the same solvent, 

We have therefore two entirely independent methods of deter- 
mining the degree of ionization of an electrolyte in solution, 
and in many cases the results obtained by the two methods 
agree very closely. 

Ostwald's Dilution Law. Suppose we have 1 gram- 
molecule of a binary electrolyte dissolved in V litres. Let 
the degree of ionization be x. Then the active masses of the 
un-ionized portion and of the two sorts of ions will be 

1 X X n X , < , 

-^ , ^ and ^= respectively. 
Hence, by the law of mass action 

k l ~~ X - F - - 
* v "VV 


~ ~~ &' "" (1 - x)V 

In other words, if x is small (as in " weak " electrolytes) the 
degree of ionization is proportional to the square root of the 
dilution, since 1 x is practically equal to 1 in these cases. 

K is called the " dissociation constant " of the electrolyte. 
Strong electrolytes, i.e. those that have a high degree of 
ionization, do not obey Ostwald's Law. 1 

1 There is reason to believe that strong electrolytes are completely 
dissociated even at comparatively high concentrations. 


Uses of the Ionic Theory. 

(i) Heat of Neutralization of Acids. Suppose we neutralize 
caustic soda with hydrochloric acid. Then 

NaOH + HC1 = NaCl + H 2 0. 

Of the four substances represented in the equation, all are 
highly ionized with the exception of the water ; water, as we 
know from the fact that it is almost a non-conductor, is 
practically non-ionized. Writing the equation in the ionic 
way, therefore, we should have 

Na + OH' + H' + 01' = Na* + 01' + H 2 0. 

The only change that has taken place, in fact, is the com- 
bination of hydrogen ions and hydroxyl ions to form un- 
ionized water. 

Since most acids, bases and salts are highly ionized in 
aqueous solution, it follows that the process of neutralization 
of an acid by a base in all thesje cases is nothing more than the 
formation of water from hydrogen ions and hydroxyl ions, 

H* + OH' = H 2 0. 

If, therefore, we take gram-equivalents of acids and neutralize 
them with gram- equivalents of bases the essential change is 
the formation of 18 grams of water from 1 gram of hydrogen 
ions and 17 grams of hydroxyl ions. Hence the heat of 
neutralization of an acid is merely the heat of formation of 
18 grams of water from hydrogen ions and hydroxyl ions ; we 
should therefore expect all acids to have the same heat of 
neutralization, which is roughly true. 

(ii) Solubility Product. If a solid electrolyte is in contact 
with its saturated aqueous solution, the following equilibria 

solid ^"" N undissociated molecules in solution ^ s ions. 

Suppose we have a binary electrolyte in equilibrium with its 
saturated solution. Let the active mass of the undissociated 
part in solution be c, that of the one sort of ion a, and that of 
the other sort of ion 6. Then by the law of mass action, 

a. b = K.c. 


Now the solution is, by hypothesis, saturated and in contact 
with the solid, hence the value of c must be constant. What 
will happen if we increase a or 6 ? Obviously more un- 
dissociated molecules of the electrolyte will be formed ; that 
is, c will be increased. It was, however, already at its maxi- 
mum value, therefore the excess of undissociated molecules 
cannot remain in solution And must be precipitated. 

This deduction can easily be tested by experiment. If we 
take a saturated solution of common salt and pass hydro- 
chloric acid gas into it, we shall increase the active mass, 
6, of the chlorine ions. Hence the product ab will be greater 
than it was before, more un-ionized molecules of salt will be 
formed, and, the maximum value for c now being exceeded, 
salt should be precipitated. This is what actually does 
happen. Addition of caustic soda ought theoretically to 
produce the same effect, and does so in practice. 

The maximum value of ab for a given compound (at a 
particular temperature) is called the solubility -product of the 
compound (at that temperature). 

(iii) Addition of Salts to Acids and Bases. If to a solution 
of an acid or a base we add a salt giving one of the ions of the 
acid or of the base, the degree of ionization of the latter 
compounds will be reduced, and if the original ionization was 
small it may be suppressed almost completely. 

Thus, in a solution of ammonium hydroxide we have the 

NH 4 OH ^= NH; + OH', 

in which the equilibrium lies chiefly to the left, since 
ammonium hydroxide is a " weak " base (i.e. little ionized 
in solution). Addition of ammonium chloride, which is, like 
practically all salts, highly ionized, will greatly increase the 
active mass of the ammonium ions, with the result that the 
active mass of the hydroxyl ions will be correspondingly 
reduced arid the equilibrium NH 4 OH - ^ NH 4 * + OH' 
shifted still more to the left. This is extremely important in 
analytical chemistry. At the stage of Group III in the 
Analytical Tables there may still be in the solution (among 
other metals) iron, chromium, aluminium, zinc, cobalt. 


nickel and manganese, all of which form " insoluble " 
hydroxides, that is, hydroxides with a low solubility product. 
If ammonium hydroxide is added to a solution in which the 
ions of these metals are present, the concentration of the 
hydroxyl ions will be sufficient to cause the solubility products 
of all the hydroxides to be exceeded, and precipitation will 
occur. * 

If, however, ammonium chloride is added first, followed 
by ammonium hydroxide, the ionization of the latter is sup- 
pressed so far that the concentration of hydroxyl ions in the 
solution is no longer high enough to cause precipitation of 
the hydroxides of zinc, cobalt, nickel and manganese, while 
still sufficient to precipitate the hydroxides of iron, chromium 
and aluminium. 

(iv) Solubility of Salts of " Weak " Acids in Solutions 
of Stronger Acids. A " weak " acid is an acid that ionizes 
only little in solution. Many organic acids are weak acids ; 
so are carbonic and sulphurous acids. It is often found that 
an " insoluble " salt of a weak acid will readily dissolve in a 
solution of a stronger acid. This is satisfactorily explained 
by the ionic theory, as will be made clear by an example. 

Calcium oxalate, CaC 2 4 , is " insoluble " in water, but 
soluble in dilute hydrochloric acid. Oxalic acid is a weak 
acid. Now, no " insoluble " salt is absolutely insoluble, so 
that we may assume that when calcium oxalate is in contact 
with water a little of it dissolves and forms a saturated 
solution. Part of the dissolved salt will ionize, so that we 
shall have the following equilibria 

CaC 2 4 solid ^ CaC 2 O 4 dissolved ;= Ca" + C 2 4 ". 
The solubility product is, however, very low. If we now add 
hydrochloric acid we shall have present in solution hydrogen 
ions, chlorine ions, calcium ions and " oxalate " ions. Oxalic 
acid, however, is a weak acid, and as soon as hydrogen ions 
meet with oxalate ions they combine to form un-ionized oxalic 
acid (the ionization of which is suppressed still further by the 
excess of hydrogen ions present. See Section iii). Hence, 
oxalate ions are removed from the solution and thus the 
equilibrium Ca" + C a O 4 " ~ ^ CaC 2 O 4 is upset, and mors 


calcium oxalate will ionize to try to restore it. This attempt 
in turn upsets the equilibrium 

CaC 2 4 ;=r CaC a 4> 

Dissolved. Solid. 

and therefore more calcium oxalate will dissolve. If suffi- 
cient hydrochloric acid is present the process will continue 
until all the calcium oxalate has passed into solution. 

(v) Hydrolysis. the salts of weak acids with strong bases 
show an alkaline reaction in solution and those of weak bases 
with strong acids show an acid reaction. This phenomenon 
is easily understood in the light of the ionic theory. Although 
water is a very poor conductor, it nevertheless is slightly 
ionized into H* and OH' and there is therefore an equilibrium 

H 2 ;== H- + OH', 

lying very largely to the left. Now the concentration of 
hydrogen ions formed from a weak acid, and of hydroxyl ions 
formed from a weak base, may be less than the concentration 
of those ions in pure water. If so, certain changes occur, 
which will be most readily understood by considering a 
definite example. Let us take the case of sodium carbonate, 
the salt of a strong base (NaOH) with a weak acid (H 2 C0 8 , 
carbonic acid). In solution this salt will ionize into Na*, Na* 
and CO 3 ". C0 3 " ions, however, immediately combine with 
the hydrogen ions of the water to form un-ionized carbonic 
acid, H 2 CO 3 ; this disturbs the equilibrium H*+OH'^=H 2 O, 
and more water ionizes. The hydrogen ions formed, how- 
ever, are immediately seized upon by C0 3 " ions and so the 
process goes on until practically all the C0 3 " ions have 
become converted into un-ionized H 2 CO 3 . But for every 
hydrogen ion removed in this way, there remains in solution 
a hydroxyl ion, and the accumulation of these hydroxyl ions 
in the solution gives the liquid an alkaline reaction. 

(vi) Indicators. Substances like litmus, phenolphthalein 
and methyl orange, which change colour with acids and 
alkalis, are themselves either weak acids or weak bases, 
which on ionization produce ions differing in colour from the 
undissociated molecules. Thus methyl orange acts as a weak 
base, and in solution gives colourless hydroxyl ions and red 



cations. The undissociated substance is yellow, hence when 
dissolved in water methyl orange gives an orange-red solution. 
Addition of alkalis suppresses the ionization of the indicator 
and the solution therefore turns yellow the colour of the 
undissociated molecules. Addition of acids, on the other 
hand, turns the solution red, since the hydrogen ions of the 
acid combine with hydroxyl ions of the indicator and further 
ionization of the latter occurs. The colour of the cations 
thus becomes apparent. 

Phenolphthalein behaves as a weak acid ; the undissociated 
molecule is colourless while the anion is pink. Addition of 
an acid produces no obvious effect, but addition of an alkali 
increases ionization and the pink colour of the anions appears. 

The sensitiveness of an acid indicator depends upon the 
concentration of hydrogen ions necessary to suppress its 
ionization and upon that of hydroxyl ions necessary to make 
its ionization great enough for the characteristic colour of the 
anions to appear. 

The active mass of hydrogen ions in pure water is 10~ 7 ; 
that of hydroxyl ions is of course the same. It is clear that 
in a neutral solution the concentration of hydrogen ions must 
be equal to that of hydroxyl ions, and this concentration 
must be the same as that of the ions of water, namely, 10~ 7 . 
An excess of hydrogen ions in solution renders the liquid acid, 
while an excess of hydroxyl ions makes it alkaline. 

A good indicator is therefore one that changes colour 
with a small excess of hydrogen ions or of hydroxyl ions ; 
judged by this criterion, the best indicator is undoubtedly 
litmus, which is blue when the concentration of hydrogen 
ions is 10~ 7 and red when the concentration is 10~ 6 . Methyl 
orange changes colour at a concentration of hydrogen ions of 
10~ 4 , and phenolphthalein at 10~ 9 . Methyl orange is there- 
fore obviously useless for titrating weak acids, but may be use- 
ful for weak bases ; with phenolphthalein the contrary is true. 

Hydrogen ion concentration has now become so extensively 
used in examination of the properties of liquids, not merely 
in pure chemistry but in medicine, agriculture, physiology 
and industry, that a special symbol has been introduced to 


represent it. This symbol is p R , which stands for the 
logarithm to the base 10 of the true concentration, with the 
minus sign omitted. Thus p R = 6-8 indicates a hydrogen 
ion concentration of 10~ 6 ' 8 . 

(vii) Strengths of Acids (and Bases). The strengths of acids 
(and bases) are compared by measuring their degree of 
ionization at equivalent concentrations. Thus, the degree of 
ionization in normal solutions of sulphuric, nitric and hydro- 
chloric acids is shown in the table 

N -Sulphuric acid . . . 0-51 
N- Hydrochloric acid 0-78 

N-Nitric acid 0-82 

In normal solution, therefore, sulphuric acid is less than 
two -thirds as " strong " as hydrochloric acid or nitric acid. 


1. Write an account of the ionic theory. 

2. What is van't Hoff's coefficient ? 

3. Define degree of ionization. Show how the degree of ionization 
of an electrolyte in solution may be calculated. 

4. State, deduce and interpret Ostwald's Dilution Law. 

5. Show that the ionic theory affords a satisfactory explanation of 
the fact that the heat of neutralization of strong acids is fairly constant. 

6. Calcium phosphate is insoluble in water, but dissolves in dilute 
hydrochloric acid. Explain this. 

7. Explain, in terms of the ionic theory, the precipitation of salt 
from its saturated solution on addition of hydrochloric acid gas. 

8. Discuss the use of (a) ammonium hydroxide, (6) sulphuretted 
hydrogen, in qualitative analysis. 

9. Write a short account of the theory of indicators. 

10. Calculate the degree of ionization of a solution of common salt 
given that the solution contains 0-724 gm. of salt in 100 gms. water and 
freezes at 0-44 C. The molecular depression of the freezing point 
of water is 18-6 and the molecular weight of sodium chloride is 58-5. 

11. The equivalent conductivity of a 5 per cent, solution of hydro- 
chloric acid is 281, while the conductivity at infinite dilution is 349. 
Find the degree of ionization of the acid. 

12. Find tha freezing point of a solution of acetic acid containing 
1 gram -molecular weight in 120 litres, given that the conductivity of 
the solution is 16-75, and that the conductivity at infinite dilution ia 
364. The molecular depression of the freezing point for water is 18'd 
and the molecular weight of acetic acid is 60. 

13,, Explain what is meant by the symbol p^. 



Many chemical reactions can be hastened, and many can 
be retarded, by the addition of small quantities of substances 
that are left unchanged in mass and in chemical composition 
at the end of the reaction. This phenomenon is called 
catalysis and the substances that bring it about are called 
catalysts or catalytic agents. Catalysts that increase the 
speed of a reaction are called positive catalysts ; those that 
retard an action are called negative catalysts. Positive 
catalysts are much more common than negative catalysts, 
hence " catalyst " generally means positive catalyst, just as 
we write a for + a. 

The first example of catalysis usually met with is the use 
of manganese dioxide in the preparation of oxygen from 
potassium chlorate by the action of heat. If potassium 
chlorate be heated alone, oxygen is given off at a compara- 
tively high temperature, but if a little manganese dioxide be 
mixed with the chlorate the latter yields its oxygen much 
more quickly at the same temperature, or at the same rate at 
a lower temperature. Since the rate of a chemical change is 
roughly doubled for a rise in temperature of 10, it is clear 
that the reaction has been hastened in both of the above cases. 
The manganese dioxide, however, is unchanged in weight and 
in chemical composition at the end of the reaction ; a fact 
which may easily be proved by adding a known weight of 
manganese dioxide to some potassium chlorate, heating until 



all the oxygen is driven off and then adding sufficient water 
to the residue to dissolve all the potassium chloride left. If 
the residual manganese dioxide is then filtered off, washed, 
dried, and weighed, its weight will be found equal to the 
original weight. 

Many other substances will also act as catalysts in the 
decomposition of potassium chlorate by heat, e.g. ferric oxide, 
Fe 2 s , and cupric oxide, CuO. 

Ostwald compared the action of a catalyst upon a chemical 
reaction to that of oil upon a machine. The similarity is, in 
fact, very great as we can see from the following points 

(i) The catalyst is left unchanged in weight and in chemical 
composition after the reaction : lubricating oil is not con- 
sumed by lubrication although it may be scattered and lost 
in this and other ways. 

(ii) The addition of a catalyst affects the speed of the 
reaction only, not the products ; similarly, if oil be added 
to a machine making paper bags, the machine will continue 
to make paper bags and not packets of cigarettes. 

(iii) Within certain limits, the more catalyst is added, the 
greater its effect : the same is true of oil. 

(iv) A catalyst will not start a reaction as a rule, and in no 
case unless the action is potentially possible : it is of no use 
adding oil to a machine if the necessary driving force is absent. 

(v) A catalyst may be either positive or negative : axle- 
grease will increase the speed of a railway-truck but would 
probably retard the action of a watch. 

(vi) A substance which is a catalyst in one reaction may not 
be a catalyst in another reaction in which it is used (e.g. 
manganese dioxide in (a) decomposition of potassium chlorate 
and (6) oxidation of hydrochloric acid to chlorine) ; similarly, 
the same oil which is a " catalyst " when used for lubrication 
may be converted into margarine and used for food. 

In the case of a reversible reaction the final state of equili- 
brium is not affected by use of a catalyst, although the time 
taken for the establishment of that equilibrium may be altered 
considerably. If the catalyst did affect the final equilibrium 
we could obtain a continuous supply of energy, out of nothing, 


by allowing the reaction to take place alternately in the 
presence and absence of the catalyst. As the creation of 
energy is entirely opposed to all our ideas and observations 
on nature, we must assume that the catalyst cannot affect 
the final state of equilibrium in a reversible reaction ; it 
therefore affects the velocities of the forward and reverse 
reactions to the same extent. This seems a very significant 
fact for the theory of catalysis, but it must be admitted that 
the above argument has been severely criticized, and that 
certain cases are known in which a catalyst at least appears 
to alter the final state of equilibrium in a reversible reaction. 

Many examples of catalysis are known, and the idea is of 
very respectable antiquity, for, describing the properties of 
the Elixir which (supposedly) changes the baser metals into 
gold, MARY THE COPT (the inventor of the water-bath), who 
flourished in Egypt long before the Muslim Conquest, says, 
" One dirJiam l thereof is sufficient for all which lies between the 
East and the West." The catalytic effect of yeast upon dough 
is probably the earliest recorded example of the phenomenon. 

Up to the present, no completely satisfactory theory of cata- 
lysis has been suggested. It is probable that the mechanism 
of the process may differ in different cases. The formation of 
unstable intermediate compounds of the catalyst and one or 
more of the reacting substances has been suggested, and in 
many cases of catalysis this may be the correct explanation. 
However, the whole subject is still shrouded in mystery and 
we must content ourselves at this stage with recording 
examples of catalytic actions, leaving their detailed study 
for a more advanced course. 

(i) The use of copper chloride in DEACON'S process for the 
manufacture of chlorine, p. 499. 

(ii) The use of oxides of nitrogen in the lead-chamber pro- 
cess for the manufacture of sulphuric acid, p. 469. 

(iii) Catalytic combination of sulphur dioxide and oxygen 
by means of platinized asbestos, p. 473. 

(iv) Catalytic decomposition of hydrogen peroxide by 
means of finely-divided metals, p. 433. 

1 About the weight of a sixpence. 


(v) The hydrogen ions of acids will accelerate catalytically 
the hydrolysis of esters by water, e.g. 

CH 3 .COOC 2 H 6 + H 2 0-> CH 3 .COOH + C 2 H 6 OH. 

Ethyl acetate. Water. Acetic Acid. Ethyl alcohol. 

(vi) Synthetic formation (p. 340) and oxidation (p. 356) of 

(vii) The catalytic action of ferrous sulphate on the oxida- 
tion of certain organic compounds by hydrogen peroxide ; 
thus tartaric acid is oxidized to dihydroxymaleic acid 




[See FENTON, Outlines of Chemistry, p. 180.] 

(viii) Many reactions between gases occur in the presence 
of platinum black, reduced nickel, and other finely -divided 

(ix) Many substances when absolutely dry will not react to- 
gether ; they will do so, however, in the presence of even a 
trace of water. Thus dry ammonia has no action upon dry 
hydrochloric acid gas. 

(x) Enzymes. Enzymes are certain complicated com- 
pounds of carbon which are very important catalysts in many 
organic reactions. They are present in, and can be extracted 
from, living animals and plants and play an essential part in 
the life- processes. The fermentation of sugar by means of the 
yeast plant is a catalytic action. Yeast contains two enzymes 
called invertase and zymase ; the first of these acts as a catalyst 
upon the hydrolysis of the cane-sugar 

H 2 + C 12 H 22 O n = C 6 H 12 6 + C 6 H 12 6 , 

Cane-sugar. Grape-sugar. Fruit-sugar. 

which results in the formation of a mixture of grape-sugar 
and fruit-sugar. (It will be noticed that both grape-sugar 
and fruit-sugar have the formula C 6 H 12 6 ; the atoms are, 


however, arranged in different ways in the molecules of the 
two substances, and grape-sugar and fruit-sugar are said to 
be isomeric with one another.) 

The mixture of grape-sugar and fruit-sugar is then cataly- 
tically split up by the zymase into alcohol and carbon dioxide, 

CH 12 6 = 2C 2 H 5 .OH + 2CO, 



In the middle of the nineteenth century GRAHAM showed 
that if a solution of salt was placed in a parchment drum 
floating in a vessel of water, the salt gradually passed through 
the membrane into the surrounding water ; glue, however, 
if made up into a solution and treated in the same way, would 
not pass through the membrane. Further investigation 
showed that most crystalline substances behaved in the same 
way as salt, while silicic acid, gum, starch, albumen, caramel 
and certain other substances, behaved like glue. The latter 
bodies Graham therefore called colloids, because they were 
" like glue " (Greek, kolla, glue) ; those which readily 
passed through a membrane he called crystalloids. It 
follows that if we have a mixture of a colloid and a crystal- 
loid in solution, the two may be 
separated by floating the solution in 
a parchment drum in a vessel of pure 
water, when the crystalloid will slowly 
pass through while the colloid re- 
Fio. 31. Graham's mains. This process Graham called 

Dial ^ zer - dialysis. 

More recent work has shown that this classification of sub- 
stances into colloids and crystalloids is unsound ; any sub- 
stance may be obtained as a colloid under suitable conditions, 
and it is therefore preferable to talk of a colloidal state of matter, 
into which some substances pass more readily than others. 

In a colloidal solution of a solid in a liquid, the particles of 
the solid are so fine that they will not settle to the bottom, 
and yet are not so fine that the substance can be regarded 
as in a " true " solution. In practically all colloidal solutions 



the particles of the colloidal substance can be observed either 
directly or indirectly by means of the microscope or ultra- 

A colloidal solution is usually easily coagulated, that is, the 
colloid is precipitated from the solution. Often mere rise of 
temperature is sufficient to cause coagulation ; in other cases 
addition of an electrolyte will bring it about, the tervalent 
cations of iron, chromium and aluminium being especially 
active in this respect hence the use of alum to stop bleeding. 
Blood is a colloidal solution of an albuminoid substance and 
is coagulated to a clot by the tervalent aluminium ion ; the 
clot closes the ends of the capillary blood-vessels and thus the 
flow of blood is arrested. 

FIG. 32. Bredig's Method for obtaining Colloidal Metals. 
A, A. Electrodes of the metal. B. Water. C. Ice. 

On the other hand, addition of a more stable colloid to 
a colloidal solution may often make the latter much more 
resistant to coagulating agents ; this phenomenon is called 
protection. Thus a colloidal solution of graphite is rendered 
very stable by the addition of colloidal tannin (see p. 286). 

Finally, it should be noted that the particles of colloids are 
usually negatively charged, although in a few cases they are 
positively charged. The liquid in which the particles are 
suspended has an opposite charge to that on the particles 

Colloidal metals are conveniently obtained by BREDIO'S 
method, which consists in striking the electric arc between 


two poles of the metal beneath the surface of extremely pure 

In the form of a solid colloidal solution, colloidal gold 
(Purple of Cassius, p. 237) has been known since 1685, while 
a rudimentary variety was employed by the ancient Assyrians 
in the manufacture of artificial red coral. A colloidal solution 
of gold in water has recently been used in medicine ; it will 
be interesting to know whether the extravagant claims of the 
alchemists for the healing qualities of " potable gold " were 
justified ! 

The study of colloids finds many applications in industry 
and various arts and crafts. Thus dyeing, photography, the 
rubber industry, the purification of drinking-water and the 
manufacture of artificial silk all present problems in colloid 
chemistry, investigation of which has led to important and 
beneficial advances. In biology and biochemistry research 
into the nature of such colloidal substances as protoplasm and 
albumen has already thrown much light upon the mechanism 
of vital processes while even ice-cream manufacturers add 
protective colloids such as gelatin to improve the " feel " of 
their product in the mouth ! 


The Phase Rule deals with the equilibria between the 
various parts of a heterogeneous system, such as the system 
ice, water and water- vapour, or calcium carbonate, calcium 
oxide and carbon dioxide, in which each part is homogeneous 
in itself but in a different physical state from other parts. 
The homogeneous parts of the system are called phases, and 
the least number of substances from which all the phases of 
the system can be made is called the number 'of components 
of the system. Thus, in the system ice, water and vapour, 
there are three phases and one component ; in the system 
calcium carbonate, calcium oxide and carbon dioxide there 
are three phases and two components, since all three phases 
may be made from carbon dioxide and calcium oxide. 

The equilibria between the phases of a heterogeneous 
system may be affected by temperature, pressure and con- 


cent ration, and the number of these factors that must be 
arbitrarily fixed in order that the system under consideration 
shall be in a definite state of equilibrium is called the number 
of degrees of freedom of the system. 

The Phase Rule may then be stated as follows 
The number of phases + the number of degrees of freedom 
= the number of components -f 2, 
or P + F = C + 2. 

Consideration of the following examples will render the 
simple application of the Phase Rule easily intelligible ; for 
more advanced work the standard textbooks on the subject 
must be consulted the student is recommended to read the 
excellent and fascinating book by Professor FINDLAY (The 
Phase Rule, Longmans & Co.). 

(i) Water and Water- vapour. Here we have 2 phases 
and 1 component, hence by the phase rule the number of 
degrees of freedom is 1, i.e. the system is univariant. This 
means that if we fix the temperature at which the water and 
water-vapour are to exist together in equilibrium, then the 
pressure will adjust itself to a definite value ; or if we wish 
to have water and water- vapour in equilibrium with one 
another, at a certain pressure of the vapour, then the tempera- 
ture will have to be adjusted to a definite value. 

(ii) Ice, Water and Water-vapour. In this system we 
have three phases and 1 component ; hence the number of 
degrees of freedom = 0, or the system is non- variant. This 
means that we cannot fix arbitrarily any of the factors govern- 
ing the equilibrium ; ice, water and water-vapour can exist 
together in equilibrium only at one particular temperature 
and pressure, dependent entirely on the system and not 
subject to any arbitrary decision on our part. 

(iii) Calcium Carbonate, Calcium Oxide and Carbon 
Dioxide. There are here 3 phases and 2 components; 
the system is therefore univariant. We can fix arbitrarily 
the temperature at which we wish the system to exist in 
equilibrium, but if we do, the pressure will automatically 
adjust itself. In other words, at a given temperature, 


calcium carbonate will give a definite pressure of carbon 

(iv) Iron, Iron Oxide, Steam and Hydrogen. The 

steam and hydrogen form a homogeneous mixture and 
thus constitute one phase ; the total number of phases is 
therefore 3. The number of components is 3, since we could 
make all the phases from iron, hydrogen and oxygen, hence 
the system is bivariant. Now the factors concerned are the 
temperature, the total pressure of the gaseous system, and its^ 
concentration or volumetric composition. If two of these 
factors are arbitrarily fixed the third will adjust itself. At 
a constant pressure, therefore, the ratio by volume of hydrogen 
to steam will be constant at a constant temperature. Of 
course, if liquid water is present in the system as well, the 
system will then be univariant. 


1. Give a general account of catalysis, with examples. 

2. Write a short essay on colloids. 

3. State and explain the Phase Rule, and in terms of this law discuss 
the following equilibria 

(i) CuS0 4 .5H 2 O ^ CuS0 4 + 5H 2 O. 
Solid. Solid. Gas. 

(ii) Ice ^ water. 
(iii) Rhombic sulphur v^ monoclinio sulphur. 


One of, the strongest instincts of mankind is to classify, 
and the chemical elements form a fascinating object for the 
exercise of this instinct. As soon as a sufficiently large 
number of elements had been isolated and described, attempts 
at a classification of them began to appear and have con- 
tinued ever since. The distinction between metals and non- 
metals was obvious and was made in remote antiquity ; it is 
convenient still, although there is no sharp line of demarcation 
between the two groups, and several elements are known 
that possess both metallic and non-metallic properties. 
These may be called the metalloids. The chief properties 
characteristic of mfctals are as follow 
(i) High density and melting-point, 
(ii) High conductivity for heat, and for electricity without 


(iii) Metallic lustre and capability of taking a high polish, 
(iv) Malleability, ductility, and great tensile strength, 
(v) Metallic oxides are usually basic. 
(vi) Metallic chlorides are generally not hydrolysed by 

(vii) Metallic hydrides are rarely formed and in any case 

are unstable. 

The main characteristics of non-metals are : 
(i) Low density. 

(ii) Poor conductivity for heat and electricity, 
(iii) Lack of metallic lustre ; will not take a polish. 



(iv) They are not malleable or ductile and their tensile 

strength is small. 

(v) Non-metallic oxides are usually acidic, 
(vi) Non-metallic chlorides are often hydrolysed by water. 

(vii) Non-metallic hydrides are common and are usually 
very stable. 

To each of these characteristics, however, both of metals 
and of non-metals, exceptions are numerous, and in deciding 
whether a particular element is a metal or a non-metal its 
general properties have to be considered as a whole. There 
is no one unfailing criterion for distinguishing between a 
metal and a non-metal, and even the general consideration 
of all the properties of an element may still leave us unde- 
cided as to its metallic or non- metallic nature. Arsenic, 
for example, possesses certain well-marked metallic proper- 
ties and other, equally well-marked, non-metallic properties. 
It is therefore said to be a metalloid. 

These facts serve to show us that while the distinction 
between metals and non-metals may be convenient in prac- 
tice it cannot be applied too rigidly. The student will find 
it a useful exercise to make a list of " non-metallic " properties 
shown by certain metals and of " metallic " properties shown 
by certain non-metals ; he may also act as judge, jury and 
advocate in deciding the claims of arsenic, antimony and 
bismuth to be considered metals and of boron to be considered 
a non-metal. 

More systematic groupings of the elements were made by 
DOBEREINER (1829) and by PETTENKOFER (1850). Dobereiner 
showed that many chemically related elements formed well- 
marked groups of three (" Dobereiner's Triads "), the atomic 
weight of the middle member of each group being approxi- 
mately the mean of the atomic weights of the other two. 
Thus the atomic weight of bromine (80) is roughly the mean 
of 35-5 and 127, the atomic weights of chlorine and iodine 
respectively. Calcium (40), strontium (87) and barium (137) 
form another such group. 

Greater progress was scarcely to be expected at the time 
owing to the uncertainty as to the atomic weights of many 


elements. However, when CANNIZZARO, in the middle of the 
nineteenth century, drew the attention of chemists to the 
great value of Avogadro's Hypothesis for deciding between 
rival values for atomic weights, and the latter were therefore 
first definitely fixed, interesting relationships became obvious 
almost at once. 

A great advance was made in 1863 by NEWLANDS, who 
pointed out that when the elements were arranged in order of 
their atomic weights, the eighth element resembled the first, 
fifteenth, etc., the ninth resembled the second, sixteenth, etc., 
and so on. Each element, in fact, more or less closely re- 
sembled the elements which were seven, or some multiple of 
seven, places before it or after it, thus recalling the arrange- 
ment of octaves in music. Newlands therefore called his dis- 
covery the Law of Octaves. Here again, however, there were 
many discrepancies, and the reception accorded to Newlands' 
work by the chemical world was the reverse of enthusiastic. 

Shortly afterwards, the Russian chemist MENDEK&EFF 
(1869) published an important paper on the classification of 
the elements, in which he described an arrangement of the 
elements that has since become famous as the Periodic 
System. Like Newlands, Mendeleeff arranged the elements 
in the order of their atomic weights, starting from the lowest, 
and called attention to the fact that a kind of periodicity 
in their properties was thus manifested, that is, chemically 
similar elements recurred at approximately equal intervals. 
This, of course, had been shown by Newlands previously 
(although it seems that Mendeleeff was not acquainted with 
Newlands' work), but Mendeleeff carried the classification 
much further, and was able to overcome many of the diffi- 
culties that had told against the Law of Octaves. 

In his Faraday Lecture to the Chemical Society in 1889, 
Mendeleeff gave the following summary of the conclusions at 
which he had arrived : 

1 . The elements, if arranged according to their atomic 
weights, exhibit an evident periodicity of properties. 

2. Elements which are similar as regards their chemical 
properties have atomic weights which are either of nearly the 


same value (e.g., platinum, iridium, osmium), or which increase 
regularly (e.g., potassium, rubidium^ caesium). 

3. The arrangement of the elements, or of groups of ele- 
ments, in the order of their atomic weights corresponds to 
their so-called valencies, as well as, to some extent, to their 
distinctive chemical properties, as is apparent, among other 
series, in that of lithium, beryllium, barium, carbon, nitrogen, 
oxygen and iron. 

4. The elements which are the most widely diffused have 
small atomic weights. 

5. The magnitude of the atomic weight determines the 
character of the element, just as the magnitude of the mole- 
cule determines the character of a compound body. 

6. We must expect the discovery of many yet unknown 
elements, for example, elements analogous to aluminium and 
silicon, whose atomic weights would be between 65 and 75. 

7. The atomic weight of an element may sometimes be 
amended by a knowledge of those of the contiguous elements. 
Thus, the atomic weight of tellurium must lie between 123 
and 126, and cannot be 128. l 

8. Certain characteristic properties of the elements can be 
foretold from their atomic weights. 

The Periodic Table in its present form is given on page 153. 
The general arrangement is due to Mendeleeff, but many 
new elements have been discovered since his time, and these 
are included. The numbers in thick black type are the 
Atomic Numbers (see Chapter XVI). 

It will be observed that the elements fall into 7 periods 
and 9 groups. The first period contains one element only, 
hydrogen. The second and third periods contain 8 elements 
each, the fourth and fifth periods 18 elements each, the sixth 
period 32, and the seventh period at present contains 6 
elements. The first three periods are called short periods 
and the rest long periods. In each period the elements show 
a gradation in properties from the chemically neutral helium 
elements through the strongly electropositive alkali metals 
to the strongly electro -negative halogens. 

1 Here Mendel e'en 7 was wrong. See p. 157. 






OCO i 




rH r-l r-t 











5 g 


CO 1 

iA r 

CO ,H 

E 1 * 



i^ S 


<r ~ 






? s 

9 S 

!. ! 


co^i PH 




^ S 






O 00 





O ^Ji 

*? <N 





5J ,0 


!> ^1 

.O QQ 

c3 W 




H 8 






* s 

? S 

c5 |aj^ 






5 fl 







S a 

o S ** ^ 







r-l OS 

9 rH 



"*-l * 

* S 

QQ ^^ 

I^ ^ 

? 5 





Qroup n. 


Mg 24-32 

r- ^ 
9 Jg 

8 f 

Sr 87-63 

48 Cd 112-40 

is | 

* 8 

Ra 226-4 






Na 23-00 

* ^ 

w 3 3 

Eb 85-45 

47 Ag 107-88 

I. i 

S g 








00 e 


























In each group, however, the elements show a similarity in 
properties. The elements of each group whfch fall in the 
periods after the first three can be divided into two sub- 

710 20 30 40 50 60 70 80 90 100 110 120 130 140 150 WO 170 180 190 200 210 220 230 240 



71020d040 506070 

Atomic Weights 

FIG. 33. Lothar Meyer's Atomic Volume Curve. 

groups ; the elements of one sub-group more closely resemble 
those elements of the group which fall in the second and 
third periods than do those of the other sub-group. Thus, 
in Group I, potassium, rubidium (Rb) and caesium (Cs) have 


a greater resemblance to lithium and sodium than have 
copper, silver and gold ; while in Group VII bromine and 
iodine are much more similar to fluorine and chlorine than 
is manganese. 

The elements in Group VIII are called the transition 
elements ; thus iron, cobalt and nickel show a gradual transi- 
tion in properties, iron being similar to manganese and cobalt, 
cobalt to iron and nickel, and nickel to cobalt and copper. 

LOTHAK MEYER (1869) pointed out that many of the 
physical properties of the elements are also " periodic func- 
tions of the atomic weight." Thus if a curve is made by 
plotting atomic weights against atomic volumes (i.e., atomic 
weight divided by specific gravity), it is found that the curve 
shows *a periodic nature (see Fig. 33), similar elements occupy- 
ing similar positions on the curve. Specific gravity, hard- 
ness, melting-point, thermal and electric conductivity, latent 
heat of fusion, and ductility of elements are all periodic 
functions of their atomic weights, and will give similar curves 
to the Atomic Volume curve if plotted in the same way. 

The specific heats of the elements, however, are non-periodic 
(cf. Dulong and Petit 's Law, p. 54). 

The valencies of the elements very often rise from in 
Group to 4 in Group IV and fall again to 1 in Group VII, e.g., 


Valency 01 234321 

Element He Li Be B C N F 

Ne Na Mg Al Si P S Cl 
Or, since the elements towards the right of the table often 
show two valencies, we may get a constant rise in valency 
from Group to Group VII, e.g., 


Valency \ 

(shown by [- Li 2 BeO B 2 8 C0 2 N 2 6 S0 3 C1 2 0, 
compound) / 
Element He Li Be B C N S Cl 


Gaps in the Table. When Mendeleeff first arranged his 
Periodic Table he had to leave 2 gaps between zinc and 
arsenic in period 4 and another between calcium and titanium, 
as no elements were then known that could fill them. 
Mendeleeff prophesied, however, that these three elements, 
which he called " efca-aluminium," " e&a-silicon," and 
" e&a-boron," would be discovered if search were made for 
them, and foretold their chief properties by a consideration 
of the properties of the neighbouring elements already known. 
A few years afterwards three elements were discovered which 
were found to have almost exactly the properties foretold by 
Mendeleeff. These elements were gallium (eka-aluminium), 
germanium (eka-silicon), and scandium (eka-boron). Thus 
what had at first appeared a weakness in Mendeleeff' s system 
was later shown to be a brilliant vindication of it. 

The Periodic System proved useful in another direction, 
namely, the correction of the atomic weights of certain elements. 
WINKLER had shown that the equivalent weight of indium 
(In) is 38, and the atomic weight was considered to be twice 
this, i.e., 76. There was, however, no place in the system for 
an element of atomic weight 76 having the properties of 
indium, and Mendeleeff therefore suggested that the valency 
of the element was probably 3 and the atomic weight 38 X 3, 
or 114. This would make indium fall into the (at that 
time) vacant space in Group III between cadmium and tin. 
Further research on indium compounds showed that Men- 
deleeff was right. The atomic weights of beryllium, uranium 
and gold were similarly corrected. 

Defects of the System. Argon has a higher atomic 
weight than potassium, cobalt than nickel, and tellurium than 
iodine. If these elements are placed in the order of their 
atomic weights they are obviously in their wrong positions ; 
argon, for example, would be with the alkali metals and 
potassium with the inactive gases ! These difficulties puzzled 
chemists for a great many years until it was discovered that, 
after all, atomic weights are not really fundamental charac- 
teristics of the elements, and that many elements exist in 
two or more forms of different atomic weights (see isotopes, 



Chapter XVI). A more fundamental property of an element 
is its atomic number or the number of resultant positive 
charges on the nucleus of the atom. If the elements are 
arranged in order of their atomic numbers the above dis- 
crepancies disappear, as shown in the following table : 


Atomic Number. 

Atomic Weight. 

Argon .... 



Cobalt .... 



Nitfkel .... 






Iodine .... 



For the position of hydrogen in the Periodic System see 
page 185. The chief defect of the system is undoubtedly the 
fact that it often separates elements of a similar chemical 
nature and groups dissimilar elements together. 


1. Write an essay on the Periodic System. 

2. What are the main characteristics of (a) metals, (b) non-metals ? 

3. In what ways has the Periodic System proved useful ? 

4. What are the chief defects of the Periodic System ? Illustrate 
your answer fully, by examples. Do you think any other system of 
classification of the elements would be an improvement on the Periodic 
System f 


In the last decade of the nineteenth century two discoveries 
were made that were destined to lead to a revolution in 
scientific thought. These were the discovery of radioactivity 
(1896) by BECQUEREL and the CURIES, and that of the X-rays 
(1895) by RONTGEN. 

In 1896 Becquerel showed that uranium salts have the 
power of acting upon a photographic plate even when the 
latter is wrapped in black paper. This can easily be shown 
by wrapping a plate in a piece of black paper and putting a 
few crystals of uranium nitrate on the paper. If the crystals 
are left in position for a day or two, on developing the plate 
it will be found that the crystals have photographed them- 
selves. Compounds of thorium behave in the same way. 
Such substances were termed " radioactive," as the action 
was supposed to be caused by rays emitted by them. 

Uranium compounds occur in the mineral pitchblende, and 
when investigating this substance Madame Curie found indi- 
cations of a much more powerfully " radioactive " body in 
it. She successfully devised methods of extracting this body, 
and showed that it was a new element, which she isolated in 
the form of a mixture of its bromide with barium bromide. 
This new element was called radium. By fractional crystal- 
lization from alcohol it was found possible to separate the 
radium bromide from the barium bromide, and in 1910 
metallic radium was prepared by the electrolysis of a solution 
of radium chloride, using a mercury cathode. The radium 
liberated at the cathode dissolves in the mercury to form an 




amalgam, from which the mercury may be distilled off, leaving 
the radium as a white metal which quickly rusts in the air, 
and which, like calcium and barium, acts upon water in the 
cold with evolution of hydrogen. It melts at about 700 C. 

Radium salts will discharge an electroscope, and investi- 
gation of this property led to the discovery that radium 
gives off three different kinds of rays, called respectively the 
a-, /?-, and y-rays. The nature of 
these radiations will be discussed 

Metals, and solutions of acids, 
bases and salts in water and 
certain other solvents, conduct 
electricity, but gases under ordin- 
ary pressures are non-conductors 
unless high potentials are employed. 
If, however, the pressure is lowered, 
it is found that gases begin to 
conduct more easily, but at still 
lower pressures exceedingly high 
potentials must be employed to 
drive the discharge through. The 
phenomena of conduction are very 
characteristic. At a pressure of 
0-01 mm. a phosphorescence is 
produced on the walls of the glass 
tube opposite the cathode. The 
cause of this phenomenon was 
investigated by SIR WILLIAM 
CROOKES, who showed that the 
phosphorescence was caused by a 
stream of exceedingly minute nega- 
tively-electrified particles which he called the Cathode Rays. 

These cathode rays are deflected by electric or magnetic 
fields in exactly the way that would be expected of a stream 
of negatively -charged particles, and are capable of passing 
through thin plates of various metals. In 1895 Rontgen 
showed that from the phosphorescent spot produced by 

FIG. 34. Crookes' Tube. 


allowing cathode rays to strike upon the end of the vacuum 
tube in which they were formed, another bea'm of rays was 
projected, of great penetrating power 1 These rays Rontgen 
called X-rays. 

The particles of which the cathode rays consist are called 
negative electrons or simply " electrons." Each negative 
electron has a mass of about y-gW f that of a hydrogen atom, 
and carries a charge equal (but opposite in sign) to that carried 
by a hydrogen ion. 

It has been shown that X-rays are similar to light vibra- 
tions except that their wave-lengths are very much smaller ; 
they can be diffracted and polarized by suitable means. The 
beam of X-rays produced from an ordinary X-ray tube con- 
sists of a mixture of rays of different wave-length, in the 
same way that white light consists of a mixture of light rays 
of different wave-lengths. For the general properties of 
X-rays, textbooks of Physics should be consulted ; it is 
sufficient for our purpose to note here a very important 
feature, namely, that every element is capable of emitting 
X-rays of wave-lengths peculiar to itself, if stimulated in an 
appropriate way. Such a way is to allow X-rays of a shorter 
wave-length to strike the substance, when the latter at once 
gives off its characteristic radiation. Now, just as the 
ordinary spectrum of an element is mapped and measured 
by means of a spectrometer, so it is possible to map and 
measure the X-ray spectrum of a substance by means of an 
instrument called the X-ray spectrometer. 

To understand how this works it is necessary to know the 
principle of an instrument called the diffraction grating. If 
ordinary white light is passed through a prism it is split up 
into light of various wave-lengths, and a spectrum may be 
produced. This analysis of light may also be brought about by 
another arrangement called the diffraction grating, which 
consists of a large number of very fine parallel lines accurately 
drawn upon a plane sheet of glass in such a way that the 
spacer separating the lines are all equal. Light which falls 
on this grating is " diffracted " or bent out of its normal 
path through an angle that is constant for a given wave- 


length of light but that differs for different wave-lengths, 
so that the grating " sorts out " the light into a spectrum. 
If the width of the space between two lines of the grating 
is known, it is possible to calculate the wave-length of any 
line in the spectrum, and it is in this way that the wave- 
lengths of rays of light are measured. 

Now X-rays are of the same nature as light-rays, but 
the wave-lengths of light-rays are several thousand times 
greater than those of the X-rays, and therefore the ordinary 
diffraction gratings are much too coarse to be of any use for 
the purpose of forming an X-ray spectrum and measuring 
the wave-lengths of the various lines. However, in 1912 
LAUB suggested that the atoms in a crystal might serve as 
the lines of a diffraction grating, and the spaces between two 
consecutive parallel planes of them as the spaces of the grating. 
If this is so, then a crystal forms a natural diffraction grating 
which should apparently be of suitable dimensions for giving 
an X-ray spectrum. Upon investigation this was found to 
be the case. When X-rays fall on a crystal they are diffracted 
in exactly the same way as light is by an ordinary diffraction 
grating. Hence, to measure the wave-length of X-rays all 
we need to know is the distance between the planes of atoms 
of a particular crystal. Fortunately it has been found 
possible to calculate this distance, and therefore to find 
the actual wave-length of any X-ray. 

The X-ray spectrometer makes use of the fact that a 
crystal will act as a diffraction grating for X-rays. The 
X-rays to be examined are passed through a slit in a sheet 
of lead and then through a second slit which serves to cut 
off any scattered radiations. The pencil of rays then impinges 
on and is diffracted from a crystal fixed by means of a piece 
of wax on a horizontal arm that can revolve on a vertical 
axis over a graduated circle. After diffraction from the 
crystal the X-rays are made to pass through a third slit into 
a tube containing a gas which is easily " ionized " (or made 
to conduct) by the rays ; sulphur dioxide is commonly used for 
the purpose. In this " ionization chamber " is an electrode 
(placed in such a position that the X-rays entering the chamber 



do not strike it) connected to an electroscope. T ne ionization 
chamber is mounted on a horizontal arm which can revolve 
around the same axis as that on which the crystal is mounted. 
To conduct the experiment, the X-rays are diffracted from 
the crystal and the ionization chamber turned until an X-ray 

FIG. 35. Plan of X-ray Spectrometer. 
Q B D. Path of raya. C. Crystal. 

passes into it, causing the gas inside the chamber to become 
ionized ; this is indicated by the electroscope. The angle 
through which the ionization chamber has been turned is 
noted, and the latter is then moved still further until the 
next X-ray passes into it, as shown by the electroscope. 



In this way the X-ray spectrum of the substance under 
observation can be measured, and the intensity of any given 
line in the spectrum is indicated by the degree to which the 
electroscope is affected. 

The wave-length of the ray is calculated from the formula 
A = 2d sin 0, when 6 is the angle at which the X-ray has been 
refracted from the crystal, and d is the distance between the 
planes of atoms in the crystal. 

If a pencil of X-rays is passed through a crystal and then 
on to a photographic plate, spots are produced on the plate, 
arranged in a symmetri- 
cal way. These spots are 
caused by the scatter- 
ing of the X-rays by the 
atoms in the crystal, and 
by constructing space- 
models from the photo- 
graphs it has been found 
possible to determine the 
spatial arrangement of the 
atoms within the crystal. 
W. L. BRAGG have shown 
that the atoms in a 
crystal of potassium 
chloride are arranged in 

the way shown in .the following diagram, the potassium 
atoms being represented by black circles and the chlorine 
by white O. 

The atoms of carbon in the diamond are arranged in groups 
of six in such a way that each carbon atom is at the centre 
of the regular tetrahedron formed by the 4 atoms nearest 
it. See fig. 37. 

The method has been extended to liquids, and the shape 
and even the size of the benzene molecule have been deter- 
mined. The shape is that of a regular hexagon, of side 
0-0000000602 cm. and thickness 0-0000000119 cm. 

Radiations from Radium. Let us return now to the 

FIG. 36. Space Lattice of Potassium 


a-, /?-, and y-rays emitted by radium. It has been shown 
that the a-rays consist of positively-charged .particles of 
atomic dimensions and of atomic weight 4. Each carries 2 unit 
positive charges. The /?-rays consist of negative electrons 
moving with a very high velocity while the y-rays are X-rays 
of very short wave-lengths. These rays are produced by the 
disintegration of the radium atoms. The atomic weight of 
radium is 226 ; when 1 atom of radium gives off an a- par- 
ticle of atomic weight 4, an atom of atomic weight 222 should 
be left. This is actually the case. It has been shown that 
the a-particle is an atom of helium carrying 2 unit positive 

charges, while the 
" element " of atomic 
weight 222 has been 
isolated and is called 
" radium emanation " 
or radon. Radon itself 
is radioactive and splits 
up into helium and a 
solid substance called 
FIG. 37. the " active deposit," 

whichis still radioactive. 

This spontaneous disintegration of atoms led scientists to 
formulate hypotheses on the structure of the atom, since atoms 
were clearly no longer to be considered as indivisible. Many 
suggestions were made, but that which agreed best with 
observed facts considered the atom to consist of an exceedingly 
minute positively -charged nucleus surrounded by a number 
of negative electrons which revolve in more or less spherical 
orbits around the nucleus. Bragg showed that the a-particles 
emitted from radium could pass through thin sheets of solid 
substances, and proved that in doing so they pass not only 
through the spaces between the atoms of these substances, but 
also actually through the atoms themselves if these happen to 
be on their path. When the a-particles pass through atoms, 
most of the particles are not deflected from their rectilinear 
path, but a small number of them suffer large deflections. 
This phenomenon is explained by assuming that when an 


a- particle passes through an atom and is not deflected thereby, 
it has not gone near the nucleus but only through the outer 
regions of the atom those in which the negative electrons 
revolve in their orbits. If we compare the atom to our solar 
system we could regard the sun as the positive nucleus and 
the planets as the negative electrons ; now it is conceivable 
that a foreign sun might rush through our solar system yet 
never come anywhere near the Sun. It seems that the 
chances of an a-particle coming within close range of the 
nucleus of an atom are about equally likely. When, how- 
ever, an a-particle does happen to pass close to the nucleus 
of an atom it is violently deflected. In LORD RUTHER- 
FORD'S words, 1 " to account for these results, it was found 
necessary to assume that the atom consists of a charged 
massive nucleus of dimensions very small compared with 
the ordinarily accepted magnitude of the diameter of the 
atom. This positively charged nucleus contains most of the 
mass of the atom, and is surrounded at a distance by a dis- 
tribution of negative electrons equal in number to the result- 
ant positive charge on the nucleus. Under these conditions, 
a very intense electric field exists close to the nucleus, and 
the large deflection of the a-particle in an encounter with a 
single atom happens when the particle passes close to the 
nucleus. Assuming that the electric forces between the 
a-particle and the nucleus varied according to an inverse 
square law in the region close to the nucleus, [Rutherford] 
worked out the relations connecting the number of a-particles 
scattered through any angle with the charge in the nucleus 
and the energy of the a-particle. Under the central field 
of force, the a-particle describes a hyperbolic orbit round 
the nucleus, and the magnitude of the deflection depends on 
the closeness of approach to the nucleus. From the 
data of scattering of a-particles then available, it was deduced 
that the resultant charge on the nucleus was about J Ae, 
where A is the atomic weight and e the fundamental 
unit of charge [i.e., is equal in magnitude to the charge carried 
by a single negative electron], . . . 

1 Proc. Roy. Soc., xcvii, p. 374 (1920), 


" Since the atom is electrically neutral, the number 01 
external [negative] electrons surrounding the nucleus must 
be equal to the number of units of resultant charge on the 
nucleus. It should be noted that, from consideration of the 
scattering of X-rays by light elements, Barkla had shown, 
in 1911, that the number of electrons was equal to about 
half the atomic weight. . . . 

" Two entirely different methods had thus given similar 
results with regard to the number of external electrons in 
the atom, but the scattering of a-rays had shown in addition 
that the positive charge must be concentrated on a massive 
nucleus of small dimensions. It was suggested by van den 
Broek that the scattering of a- particles was not inconsistent 
with the possibility that the charge on the nucleus was equal 
to the atomic number l of the atom, i.e. to the number of 
the atom when arranged in order of increasing atomic 
weight, " taking Hydrogen as 1, Helium as 2, Lithium as 3, 
and so on. 

It will be convenient here to consider the results of an 
independent line of research carried out by MOSELEY, who 
investigated the X-ray spectra of various elements by means 
of the X-ray spectrometer. He found that the X-ray spectra 
obtained in this way show two strong lines for each element, 
accompanied by a number of weaker lines (see Fig. 38). 
Of the two strong lines, one is stronger than the other and 
is called the a-line, while the weaker is called the /Mine. It 
has been shown that if v is the frequency (i.e. number of 
vibrations per second) of the a-line, and N the atomic number 
of the element, then 

t> = J (N l) 2 Xa constant. 

This constant is called Rydberg's constant and has the value 
109 677. If, therefore, we know the frequency of the a-line 
of the X-ray spectrum of an element, we can calculate the 
position that it ought to occupy in the Periodic Table, that 
is v its Atomic Number. 

1 Italics mine. E. J. H. 



This important discovery made it possible for the first 
time to call the roll of the chemical elements and to deter- 
mine how many there were and how many remained to be 
discovered. There are between hydrogen and uranium 92 

FlQ. 38. X-ray Spectra. 

possible elements, of which only two (1931) remain to be 
found namely, the two heaviest analogues of iodine and 
caesium respectively. 

Moseley's work, in fact, showed that the " properties of 
an atom were denned by a number which varied by unity 


in successive atoms. This gives a new method of regarding 
the periodic classification of the elements, for* the atomic 
number, or its equivalent the nuclear charge, is of more 
fundamental importance than its atomic weight." 1 Most 
of the physical and chemical properties of an atom depend 
upon the number and arrangement of the negative electrons 
in the atom, and these will clearly depend upon the charge 
on the nucleus. In other words, the actual mass of the atom 
is of secondary importance. 

Hence we are led to the conclusion that " it is quite possible 
to imagine the existence of elements of almost identical 
physical and chemical properties, but which differ from one 
another in mass, for, provided the resultant nuclear charge 
is the same, a number of possible stable modes of combination 
of the different units which make up a complex nucleus may 
be possible." 2 In other words, we may get atoms which 
are chemically indistinguishable and yet of different atomic 
weights ! Are we to regard such atoms as atoms of different 
elements, or as atoms of the same element ? According to 
Dalton, all the atoms of the same element have the same 
atomic weight ; therefore from this point of view atoms 
which, are chemically identical but which have different 
atomic weights belong to different elements. On the other 
hand, chemical considerations would lead us to regard atoms 
that are chemically identical as atoms of the same element. 
SODDY gave the name isotopes or isotopic elements to those 
elements which fall into the same place in the periodic 
system, and are chemically identical, but have different 
atomic weights. 

We have already seen that when an a-particle (or helium 
atom carrying two positive charges) is expelled from a radium 
atom, the product (radon) is an element which falls into 

1 Rutherford, loc. cit. 

* Rutherford, loc cit. N.B. The idea of isotopes was by this time 
(1920) well understood, having been elaborated by chemical work on 
the disintegration of radium and its congeners. In order to make 
the subject clearer the historical sequence has not been strictly ad- 
hered to. 


Group O of the periodic system, or two columns to the left of 
that in which the parent radium atom is placed. Study of 
other radioactive products has shown that this is a general 
phenomenon expulsion of an a-particle from the atom of 
an element in Group N results in the formation of an atom 
of an element which falls into Group N-2 and which has 
an atomic weight differing by 4 units from that of the parent 
atom. Further investigation has produced evidence to show 
that when one ^S-particle is expelled from the atom (probably 
from the nucleus), an atom is formed which is that of an element 
which falls into a column one to the right of that in which the 
parent element is placed, but of the same atomic weight. 
" Each of the successive places in the periodic table thus 
corresponds with unit difference of charge in the constitution 
of the atom " which is the conclusion previously arrived 
at by van den Broek. We see, too, that there is, in addition 
to the existence of isotopes, a possibility of the existence of 
different elements with the same atomic weight : these are 
called isobar ic heterotopes. Elements that differ in chemical 
properties and also in atomic weight are called heterobaric 
heterotopes. All heterotopes are separable by chemical 

The existence of isotopes suggested above is rendered still 
more probable by the following considerations. Suppose an 
atom loses an a-particle by radioactive change. We have 
seen that an atom will be formed of atomic weight 4 units 
less, and belonging to an element 2 columns to the left ift 
the periodic table. Suppose now this daughter-atom loses 
2 ^-particles. It will have moved two places to the right in 
the table and will therefore have reached the position from 
which it set out, with no further change in atomic weight. 
We should now have two atoms differing by 4 units in atomic 
weight, but absolutely identical in chemical properties, that is, 
they are isotopic elements, or isotopic forms of the same 
element with different atomic weights. Fig. 39 will make 
this clear. 

It will be seen that atoms A and B occupy the same position 
in the table, and are chemically identical ; but they differ 




in atomic weight by 4 units : they are isotopes. C and D 
are isobaric heterotopes. 

The first case in which these views were tested experi- 
mentally was that of lead. It had been proved that the end- 
products of the radioactive disintegrations of thorium and 
of uranium both fell into the place in the periodic table 
occupied by lead, but a consideration of the intermediate 
stages led to the conclusion that the lead derived from uranium 
should have an atomic weight of 206, while that from thorium 

FIG. 39. 

should have an atomic weight of 208. Now uranium minerals 
often contain small quantities of lead and it is reasonable to 
suppose that this lead has been derived from uranium by 
radioactive changes ; similarly, the lead found in thorium 
minerals has probably been derived from thorium. Lead was 
extracted from both these sources and the atomic weights of 
the specimens carefully determined by chemists skilled in 
atomic weight determinations. It was found that the lead 
from uranium minerals had an atomic weight of 206*05 and 
that from thorium minerals 207-9 ! Thus the theory was 
triumphantly justified. Ordinary lead, of atomic weight 


207-2, is a mixture of these isotopes with a third isotope, 
A.W. 207, in the appropriate proportion. The 206-05 lead 
and the 207-9 lead were proved to be chemically identical a 
predicted by the theory. 

Since this date (1914), much more work has been done on 
isotopes of common elements, largely by means of SIB J. J. 
THOMSON'S positive-ray method of gas-analysis. Some account 
of this will first be given. 

If holes are cut in the cathode of a Crookes' tube, such as- 
that used for obtaining the cathode rays, it is found that a 
stream of rays travels from the cathode in the opposite direc- 
tion to that of the cathode rays. By observation of the effect 
produced on these rays by electric and magnetic fields, they 
were found to consist of positively charged atoms or molecules 
of the gas which is present in the tube under very low pressure. 
Sir J. J. Thomson showed that the magnitude of the deviation 
of these particles caused by the application of a magnetic or 
an electric field was proportional to the ratio of the mass of 

the particle and its charge, or . By using as cathode a 


fine aluminium rod with a capillary 

copper tube running through it from 

end to end, it was possible to get 

a fine pencil of the rays passing 

through the cathode and striking 

a photographic plate at a point 

directly opposite (A) (Fig. 40). If 

now a magnetic field be applied, the 

particle will be deflected and will 

strike the plate say at a point B. FlQ 

If the magnetic field be replaced by 

a suitable electric field the particle will be again deflected 

in a direction at right angles to that produced by the" 

magnetic field, and will strike the plate say at C. Under 

the combined influence of both fields directly the particle will 

move in a resultant deflection and may strike the plate say 

at point D. Now in a stream of identical particles, not all 

will be travelling with the same speed, and therefore they will 



not all strike the plate at the same spot. It can, however, be 
proved by a simple calculation that if the deflection caused 
by the electric field is x and that by the magnetic field y> then 

I/ 2 6 

is constant and equal to for all identical particles, 
^ r 


whatever their 

speed. = a 

constant is of course the 

Fio. 41. Typical Positive Ray 

equation of a parabola, and there- 
fore the photographic plate will 
on development show a para- 
bolic streak caused by a stream 
of particles all of one kind. These 
parabolas will, however, be differ- 
ent for particles of different sub- 
stances. Thus, if one particle is 
a hydrogen atom and another an 
atom of argon, the values of the 

7/ 2 

expression will be different 

in the two cases, and if we have 
a stream of hydrogen atoms 
mixed with argon atoms we shall 
get two parabolas, one caused by 
the hydrogen and one by the 
argon. If the mass of the one 
kind of particle is m and that 
of another w', suitable measure- 

ment of the parabolas will enable 

us to get the ratio and hence if we know m we can cal- 

culate ra' (assuming the charge carried by each particle to 
be the same). The advantages of this method of analysis were 
well summed up by STEWART, 1 who said (1920) : 

" In the first place, it permits the analysis of extremely 
small quantities of a gas-mixture. Thus a quantity of helium 
which does not exceed 4 X 10" 6 c.c. can be detected in a o.c. 
of air. 

1 Recent Advances in Physical and Inorganic Chemistry. 



" Secondly, it carries our knowledge further than spectrum 
analysis can do. If we have only a trace of an element mixed 
with large quantities of other gases, the spectrum of the 
mixture may fail to reveal the presence of the trace owing to 
its characteristic lines being swamped by the spectra of its 
companions ; whereas in the case of the positive ray method, 
each constituent is sifted out from the others. 

" Thirdly, in the case of a new element, an examination of 
the spectrum tells us only that new lines are present ; but 
with positive ray analysis it may be possible to go much 
further. For example, if we find two parabolas characteristic 

Fia. 42. Positive Ray Apparatus. 

A. Discharge flask. B. Cathode. D. Anode. E. Gas entry. P. Gas exit. 
P, P*. Analyser. U. Photographic plate. 

of the new element, one must arise from the atom and the 
other from the diatomic molecule. If the substance is 
monatomic there will be only one parabola, or, if there be 
more than one produced (owing to the atom taking up more 
than a single charge), we can detect the nature of the charged 

" These facts prove that in the positive ray method we have 
gained a new and formidable weapon for the chemical 
armoury. " 

Improvements of the positive-ray method of the analysis of 
gases, in the hands of ASTON and others, have shown that 



many elements are heterogeneous, that is, the " element " as 
ordinarily met with is a mixture of isotopes. The following 
table, given by Aston, shows the results of the investigations 
of 20 elements : 





Number GJ 

Masses of Isotopes, in order 
of their Intensity. 















11, 10 






Nitrogen . 





Oxygen . 





Fluorine . 





Neon . 




20, 22 





28, 29, 30 






Sulphur . 




32, 33, 34 

Chlorine . 




















Bromine . 




79, 81 

Krypton . 




84, 86, 82, 83, 80, 78 










129, 132, 131, 134, 136, 128, 

130, 126, 124 

Mercury . 




202, 200, 199, 198, 201, 204 

It will be noticed that in every case the atomic mass of each 
isotope is a whole number (0 = 16-00). 1 

To give an explanation of this, it has been suggested that 
the nuclei of other atoms are composed of hydrogen nuclei 
and helium nuclei. This theory has received support from 
work of Lord Rutherford, who has shown that by bombard- 
ing nitrogen atoms with a-particles it is possible to obtain 

1 Except in that of hydrogen, H 1-008, 


disintegration of a few of the nitrogen atoms, one of the 
products of the disintegration being swiftly-moving positively 
charged hydrogen atoms. 

Within the last few years, the electronic structure of the 
atoms of many elements has been worked out to a high 
degree of probability. Thus the hydrogen atom is considered 
to be composed of one unit positive charge, or proton, acting 
as nucleus, round which revolves one electron describing a 
circular orbit. The helium atom consists of a nucleus of 
4 protons and 2 electrons, round which revolve 2 electrons, 
i.e. the nucleus of helium as of all other elements except 
hydrogen consists of both protons and electrons the number 
of " orbital " or revolving electrons being equal to the net or 
resultant charge on the nucleus. 

The orbital electrons are believed to be arranged in con- 
centric groups or " shells " in all cases where there are more 
than 2 in the atom, i.e. in all atoms except those of hydrogen 
or helium. Putting the matter another way, we may say that 
the orbital electrons fall into groups, the orbits of those in 
the same group being of the same radius, but of greater or 
smaller radius than those of inner or outer groups. The sheila 
become " complete " when they contain 2n 2 electrons, where 
n = 1, 2, 3 or 4. 

The numbers of electrons in the shells are shown for several 
elements in the table following, where it will be observed that 
the total number of orbital electrons is, as previously stated, 
equal to the resultant charge on the nucleus, i.e. to the atomic 

It will be noticed that the number of electrons in the outer- 
most group is equal to the normal valency of the atom, except 
in the case of the rare gases. It is, in fact, believed that 
elements show a tendency to complete an outermost group or 
sub-group consisting of 8 electrons, or they do so in combination 
with other elements by parting with electrons of their own or 
acquiring those of the atoms with which they combine. An 
element that normally has a complete outer group of 8 electron* 
therefore shows no tendency to combine with other elements, 
i.e. its valency is 0. 






Electronic Groupi 














Lithium . 








Boron . 




Fluorine . 




Neon . 



















Argon . 























Zinc . . . 






Gallium . 






Krypton . 




















Silver . 







Cadmium . 






























Gold . . . 








Mercury . 

















The electronic groups in the figure, except Group I, may 
be still further subdivided. 

When sodium, with one valency electron, combines with 
fluorine, with 7 valency electrons, we assume that the sodium 
atom parts with its valency electron and is thus left with a 
complete outer ring of 8 electrons. The fluorine atom takes 
up the valency electron of the sodium, and thus completes its 
outer ring of 8. But the sodium atom thereby becomes 
positively charged (as it was electrically neutral before) and 


the fluorine atom similarly becomes negatively charged. The 
two atoms are therefore bound together by electrostatic 

All the orbital electrons in an atom and possibly the 
nucleus as well are believed to be in constant and rapid 
motion, hence any conception of the atom that shows them 
at rest (e.g. that suggested by LEWIS and LANGMUIB) is 
necessarily inadequate. The theories of BOHR, SCHROEDINGEB, 
DIKAC and others are, however, too complicated for us to 
consider here ; they deal with the movement, spin and 
magnetic moment of the electrons and the nucleus. 

In this chapter we have seen the opening stages of what 
will surely be progress more wonderful than was witnessed 
even by the nineteenth century. Dalton's Atomic Theory is 
shown to be a rough generalization only amply sufficient for 
everyday use and likely to be so for a very long time but 
inaccurate in every particular. If we wonder at the geniuB 
of men like Dalton and Avogadro who could make measure- 
ments, by the eye of faith, of atoms and molecules, what are 
we to think of the chemists and physicists of to-day, who deal 
with particles so small that in comparison with them an 
atom is like our solar system compared with the earth ? 

What more striking example could we desire of that great 
truth of Science, so ably expressed by one of her most brilliant 
followers, JABIR IBN HAYYAN : 

" Scientists delight not in abundance of material ; they 
rejoice only in the excellence of their experimental methods " ? 


1. What do you know of (a) positive-ray analysis, (6) X-ray spectra, 
(c) atomic numbers ? 

2. Give an account of modern views on the structure of the atom. 

3. Explain fully the term isotopes. 

4. Write an essay on a Clifton boy's dictum : " Chemistry is a out- 
destroying subject." 





Group in Periodic System : unique ; Symbol : H ; Valency : 
1 ; Atomicity : 2 ; Atomic Weight : 1-008. 

History. Although hydrogen must have been obtained 
by the alchemists, it was first recognized as a definite substance 
by CAVENDISH (1766), who prepared it by the action of dilute 
sulphuric acid upon iron or zinc. He called it inflammable air t 
and considered it to be practically pure phlogiston (p. 399). 
It was called hydrogen (" water-producer ") by LAVOISIER, 
on account of the fact (discovered by Cavendish) that a mix- 
ture of it with half its volume of oxygen on explosion yields 

Occurrence. Hydrogen is one of the most widely-dis- 
tributed of the elements, not only on the earth but through- 
out the universe the atmosphere of the sun, for example, 
consists largely of incandescent hydrogen. Free hydrogen 
occurs to some extent on the earth, in volcanic and natural 
gas and hence but in the most minute traces in the atmo- 
sphere. Combined hydrogen is found in abundance. Thus 
i by weight of water consists of hydrogen, petroleum is a 
mixture of compounds of carbon and hydrogen, and hydrogen 




is an essential constituent of all living matter. The per- 
centage of hydrogen in the earth's crust is about 0-95. 

Formation. Pure hydrogen is obtained by the electrolysis 
of a dilute solution of barium hydroxide ; the hydrogen is 
evolved at the cathode and pure oxygen at the anode. 
The gases are, however, moist and must be dried by 
standing over phosphorus pentoxide (P 2 5 ). Electrolysis 
of dilute sulphuric acid yields a rather less pure hydrogen, 
and some of the oxygen is lost owing to the formation of 
persulpburic acid (H 2 S 2 8 ) and hydrogen peroxide (H 2 2 ). 

Hydrogen may also be obtained 
from water by the action of certain 
metals, which are oxidized by the 
oxygen present and the hydrogen 
thus set free. Sodium and potas- 
sium decompose water in the cold, 
forming the hydroxide of the metal, 
and hydrogen 

2Na + 2H 2 + 2NaOH + H 2 . 

The heat evolved is sufficient to 
melt the metals, which float on the 
water as molten globules, and in the 
case of potassium the temperature 
becomes high enough to cause the 

hydrogen to burn, the flame having a lilac colour owing to 
some of the potassium burning as well. The action of sodium 
and potassium upon water is not suitable as a means of pre- 
paring hydrogen in the laboratory, as the action is too violent. 
If, however, sodium is dissolved in mercury, the sodium 
amalgam so formed will decompose water more slowly, and 
as the amalgam sinks in water the hydrogen evolved may be 
easily collected by inverting a gas-jar full of water over the 
lump of amalgam in a pneumatic trough (Fig. 43). 

Calcium is heavier than water and therefore sinks in it. It 
decomposes water in the cold 

Ca + 2H 2 - Ca(OH) 2 + H t , 
forming calcium hydroxide and hydrogen. 

FIG. 43. 



Magnesium will not attack cold water but acts upon 
hot water and readily burns in steam (Fig. 44), yielding 
magnesium oxide and hydrogen 

Mg + H 2 = MgO + H 2 , 

Magnesium amalgam and aluminium amalgam both liberate 
hydrogen from water, the former in the cold and the latter on 
heating. The action of aluminium amalgam on boiling water 
has been used for the manufacture of hydrogen. 

LAVOISIER showed that hydrogen could be obtained by 
passing steam over heated iron. 

3Fe + 4H 2 O ^= Fe 3 4 + 4H t . 

FIG. 44. 

This action is reversible (p. 119). It is used commercially 
for the manufacture of hydrogen, the iron oxide formed being 
reduced to the metal again by treatment with water-gas (a 
mixture of carbon monoxide and hydrogen made by passing 
steam over red-hot coke, C + H 2 CO + H 2 ). If a catalyst 
be employed (ferrous chloride and copper), liquid water under 
pressure readily reacts with iron at a temperature of 300 
340 C., with liberation of hydrogen, which, as it is produced 
under pressure, may be passed directly into the steel cylinders 
in which it is sold, without further compression. 


Hydrogen is generally prepared in the laboratory by the 
action of dilute sulphuric acid upon zinc 

Zn + H 2 S0 4 = ZnS0 4 + H 2 . 

For this reaction to take place at an appreciable rate, it is 
necessary that impurities should be present. If pure zinc 
and pure dilute sulphuric acid are used the evolution of gas 
is extremely slow, but addition of a little copper sulphate or 
platinum chloride will hasten the reaction. 

Other metals, such as iron and magnesium, may be used, 
and hydrochloric acid may be substituted for sulphuric acid. 
Nitric acid, however, is not suitable, as it is readily reduced 
by hydrogen, and the product of the action of nitric acid on 
metals is therefore not hydrogen but a reduction product, 
or a mixture of reduction products, of the acid (see p. 357). 

Hydrogen is evolved when zinc, tin or aluminium is heated 
with caustic soda or caustic potash solution 

i. Zn + 2NaOH - Na 2 ZnO 2 + H 2 . 

Sodium zincate. 

ii. 2H 2 + 2A1 + 2NaOH - 2NaA10 2 + 3H,. 

Sodium meta-aluminate. 

iii. Sn + 4NaOH = Na 4 Sn0 4 + 2H 2 . 

Sodium stannate. 

" Silicol " process. An alloy of silicon and iron (ferrosilicon) 
is added to a concentrated solution of caustic soda, when 
sodium silicate and hydrogen are formed, roughly according to 
the equation 

Si + 2NaOH + H 2 = 2H 2 + Na 2 Si0 3 . 

In reality the reaction is rather more complicated. A mixture 
of slaked lime and caustic soda gives better results than 
caustic soda alone. From a commercial point of view the 
chief disadvantage of this method lies in the fact that the 
materials are expensive. 

" Hydrolith " process. Hydrolith is the trade name foi 
impure (90 per cent.) calcium hydride, CaH 2 . This substance 
is prepared by heating calcium in hydrogen, and on treatment 


with cold water it yields calcium hydroxide and hydrogen 
CaH 2 + 2H 2 - Ca(OH) a + 2H 2 . 

Most of the hydrogen of commerce is obtained as a by- 
product in the manufacture of sodium and caustic soda by 
electrolytic methods. A good deal, however, is prepared from 
water-gas (p. 301) by the removal of the carbon monoxide, 
e.g. by passing the gas over heated calcium carbide, when 
a mixture of calcium carbonate and calcium oxide is left, the 
hydrogen being unaffected. 

Properties. Hydrogen is a colourless gas, with no taste 
or smell. It is the lightest substance known (1 litre weighs 
0-0896 grams at N.T.P.), and is therefore used for filling 
airships, although its inflammable nature makes it a constant 
source of danger. Owing to the serious disasters that have 
occurred from this cause, the modern practice is to substitute 
helium for hydrogen where this can be done. Helium has 
about 92 per cent, of the lifting-power of hydrogen and is 
absolutely non-inflammable. It is, however, considerably 
more expensive than hydrogen. 

Liquid hydrogen was first obtained by SIR JAMES DEWAB 
in 1895. It is a colourless liquid, of density 0-07, which boils 
at 253 C. By rapidly evaporating liquid hydrogen under 
reduced pressure, solid hydrogen can be obtained, in the form 
of white crystals melting at 259 C. It was at one time 
supposed that hydrogen was the vapour of a metal, but this 
idea was disposed of by the production of solid hydrogen, 
which has no metallic properties. 

Occlusion of Hydrogen. Many metals have the power 
of absorbing hydrogen in the cold and giving it up agaia 
when heated. Palladium possesses this power in a high 
degree 1 volume of palladium wire will absorb 930 volumes 
of hydrogen at N.T.P. Cobalt powder will absorb about 
100 volumes. Graham called this phenomenon the occlusion 
of hydrogen by metals. It was at one time supposed that in 
the case of palladium a palladium hydride, Pd 2 H 1 , was formed 

1 Or 


the hydrogen probably dissolves in the palladium forming a 
solid solution. The question is, however, still urisettled. In 
any case, there seems to be a condensed layer of hydrogen on 
the surface of the occluding metal, and this probably accounts 
for the fact that a mixture of oxygen and hydrogen is ignited 
by the introduction of finely-divided platinum : the oxygen 
and hydrogen are brought into close contact at the surface 
of the metal, when reaction between them occurs, and the 
heat evolved is sufficient to ignite the rest of the mixture. 

Charcoal will absorb hydrogen and many other gases, 
especially when it is cooled in liquid air. McBAiN (1909) 
has shown that this absorption takes place in two stages : 
(i) rapidly on the surface of the charcoal, and (ii) slowly by 
diffusion into the interior of the charcoal. The former pro- 
cess he calls adsorption, the latter absorption, and the two 
together sorption. 

Reduction. When hydrogen is passed over a heated 
metallic oxide, the latter often loses its oxygen, and the metal 
is liberated 

metallic oxide + hydrogen = metal + water. 

This process is called reduction, and hydrogen is said to be 
a reducing agent. The phenomena of reduction and oxidation 
are considered in more detail on page 428. 

" Nascent " Hydrogen. Hydrogen is much more chemically 
active at the moment of its liberation in an exothermic 
chemical reaction. It is then said to be in the nascent state. 
Thus, if hydrogen is passed through ferric chloride solution, 
no change occurs, but if zinc and dilute hydrochloric acid are 
added to ferric chloride solution, the yellow colour of the 
ferric salt disappears, and ferrous chloride is formed 
FeCl 3 + H = FeCl, + HC1. 

Many similar reactions are known. It was suggested that 
the atoms of hydrogen in the nascent state have not joined 
together to form the diatomic molecules of ordinary hydro- 
gen, and that the greater chemical activity of nascent hydro- 
gen is due to this cause. However, the chemical activity 
of nascent hydrogen varies according to the reaction in which 


it is produced, and therefore cannot be due merely to the fact 
that the hydrogen is in the atomic state. Since " nascent " 
hydrogen is formed only in exothermic reactions, the 
phenomena it produces may be due to the fact that it has 
considerable but varying amounts of energy temporarily at 
its disposal. 

" Active " hydrogen has been prepared by LANGMUIB by 
heating filaments of tungsten, palladium or platinum to 
very high temperatures in hydrogen at low pressures. It is 
much more chemically active than ordinary hydrogen, and 
can react directly with oxygen, phosphorus and other sub- 
stances, at room temperature. It is supposed that this 
hydrogen is in the atomic state. Langmuir has also shown 
that by blowing hydrogen into the electric arc the diatomic 
molecules are split up into single atoms. Combustion of 
this " atomic " hydrogen yields extremely high temperatures 
considerably above that of the electric arc and atomic 
hydrogen blowpipes have found application in the welding of 
metals, etc. 

Commercial Uses of Hydrogen. Hydrogen is used for 
filling airships and balloons ; mixed with carbon monoxide 
it forms water-gas, used as a fuel ; the flame of hydrogen 
burning in oxygen is the oxyJiydrogen flame, which is produced 
by means of special blowpipes and is used for many purposes. 
SABATIEB and SENDERENS showed that many oils could be 
converted into solid fats by reducing them with hydrogen, 
using reduced, nickel (that is, nickel obtained in a fine state 
of division by reduction of nickel oxide) as a catalyst. This 
process is used in the manufacture of margarine. Hydrogen 
is also used in the preparation of quick- drying varnishes and 
in the manufacture of synthetic ammonia. It is one of the 
chief constituents of coal-gas. 

Position of Hydrogen in the Periodic System. 

Hydrogen is a unique element it has no analogues. Some 
chemists consider it to be most closely related to the alkali 
metals, others to the halogens. The former position seems 
to be the more natural, as the alkali metals are, like hydrogen, 



electropositive while the halogens are electronegative. The 
following table shows the main points to be considered 







Molecules diatomic 

Molecules mona- 

Gases (lower mem- 
Molecules diatomic 



Oxide stable 

Oxides stable 

Oxides unstable 





Hydrides unstable 
Replaceable by hy- 
drogen in salts 

E lee tronegat i ve 
Hydrides stable 
Replaceable by hy- 
drogen in organic 
compounds, but 
not in salts 


More gaps are left 
placed above Lit 

in the Periodic Sy 

stem if hydrogen is 

Of these, i, ii and vii are favourable to the view of hydro- 
gen as a halogen, the rest show its similarity to an alkali 
metal. It has recently been shown that fused lithium hydride, 
LiH (p. 191), on electrolysis yields hydrogen at the anode. 
This would support the view of those who class hydrogen with 
the halogens. 


1. Describe the preparation and properties of hydrogen. 

2. How is hydrogen made on the commercial scale ? What are its 
uses ? 

3. Write an account of the occlusion of hydrogen by metals. 

4. Write short notes on (a) nascent hydrogen, (b) active hydrogen 
6. Discuss the position of hydrogen in the Periodic Table, 

6. What do you know of the history of hydrogen ? 


Group in Periodic System : ; valency : 0. 

Very interesting accounts of the discovery of the inactive 
gases are to be found in SIR WILLIAM RAMSAY'S book, The 
Gases of the Atmosphere, and 'in M. W. TRAVERS* The Discovery 
of the Rare Gases (Arnold, 1928), which the student is strongly 
recommended to read. 

In 1785 CAVENDISH published (Phil. Trans., 75, 372, 1785) 
an account of his work on the conversion of atmospheric 
nitrogen into oxides of nitrogen by sparking with oxygen. 
He found that a bubble of gas, about y^th of the original 
volume of nitrogen taken, defied all efforts made to convert 
it into nitrogen oxides, and said, " If there is any part of the 
phlogisticated air of our atmosphere which differs from the 
rest ... we may safely conclude that it is not more than 
Tihrth part of the whole." 

For over a century Cavendish's observation was neglected, 
but in 1894 RAYLEIGH found that the density of atmospheric 
nitrogen was always slightly greater (about 0-5 per cent.) 
than that of nitrogen obtained by the decomposition of 
nitrogenous compounds. Rayleigh then repeated Caven- 
dish's experiments and found that his results were correct ; 
spectroscopic examination of the residual gas showed that it 
certainly was not nitrogen. 

Rayleigh next began experiments, in collaboration with 
SIR WILLIAM RAMSAY, to obtain the new gas in larger quan- 
tities. This was done by absorbing the oxygen of air by 



red-hot copper and the nitrogen by red-hot magnesium. In 
this way a gas was obtained of density 19-94. The new gas 
was characterized by its entire lack of chemical activity, and 
was therefore called argon (Creek, idle). Since argon forms 
no chemical compounds at all, it has no equivalent j 
the molecular weight, however, must be twice the vapour 
density (19-94) and is therefore 39-88. The atomicity of the 
molecule was found to be 1> by determining the ratio of the 
specific heats of the gas (p. 41) ; hence the atomic weight 
also is 39-88. 

Argon was afterwards prepared on a larger scale by the 
fractional distillation of liquid air, and careful examination 
of the different fractions led to the discovery of four more 
gases similar to argon helium, neon, krypton and xenon. 

SIB NORMAN LOCKYBB had many years previously (1868) 
discovered helium in the sun, by means of the spectroscope, 
and it was also found by Ramsay in 1894 in the gases evolved, 
on heating, from the mineral cleveite. More recently, helium 
has been found in large quantities (over 1 per cent.) in the 
natural gas of certain American (Texas) gas-wells, and has 
been extracted on a commercial scale for use in airships, since 
it is very light (V.D. = 2) and non-inflammable. 

The a-particles expelled from radium were shown to consist 
of charged helium atoms, and modern work on the structure 
of matter (Chap. XVI) affords strong evidence for the theory 
that the nuclei of all other elements are composed of hydrogen 
and helium nuclei. 

A sixth member of the group is radon, or radium emanation, 
a radioactive gas evolved from radium. It is unstable, and 
disintegrates fairly quickly into helium and other products ; 
this disintegration is not a chemical change in the ordinary 
sense of the term, but an actual decomposition of radon atoms. 

The inactive gases (or helium group of elements) fall 
naturally into the periodic system between the strongly 
electropositive alkali metals and the strongly electronegative 
halogens. The atomic weight of argon (39-88) is slightly greater 
than that of potassium (39-1) and argon therefore appeared 
to be an anomaly in the system until it was realized that the 



atomic numbers (p. 166) of the elements, and not their atomic 
weights, are the more fundamental characteristics. The 
atomic number of argon is 18, whereas that of potassium is 
19 ; by taking atomic numbers instead of atomic weights, 
therefore, the discrepancy disappears. 

All the inactive gases are without chemical properties of 
any kind, and their molecules are all monatomic. The chief 
physical characteristics are given in the following table : 


Weight. Number. 



Ratio of Specific 

Neon . 

4-0 2 

20-2 10 


1-5 (?) 


Krypton . 

39-88 18 
82-92 36 
130-2 54 




Radon' . 

222-6 86 




Helium proved extraordinarily difficult to liquefy ; it was 
first liquefied by ONNES in 1907. 

The following isotopes (p. 168) of the inactive gases are 
known ; partial separation has been effected in the case of 
neon : 

Helium, 1, At. wt. 4-0. 

Neon, 2, At. wts. 20 and 22. 

Argon, 2, At. wts. 40 and 36. 

Krypton, 6, At. wts. 84, 86, 82, 83, 80 and 78. 

Xenon, 7, At. wts. 128, 129, 130, 131, 132, 134 and 136. 

Neon and argon are prepared commercially for use in gas- 
filled electric lamps ; indeed, the cheapest way of obtaining 
a specimen of neon is to buy a neon lamp. Of recent years, 
neon electric discharge tubes have played an important part 
in the illuminated advertisements of our great cities ; at night 
the Boulevard des Italiens, Paris for example is diffused 
with the characteristic and beautiful orange-red neon glow. 


Neon lighthouses are also used at the principal aerodromes, 
especially to guide aeroplanes in times of fog, since the light 
rays at the red end of the spectrum are very penetrating in 
foggy conditions, as is evidenced by the colour of the sun on 
a foggy day. 


1. Give an account of the discovery of argon. 

2. Describe the general properties of the inactive gases. 

3. Discuss the position of the inactive gases in general, and of argon 
in particular, in the Periodic System. 


TYPICAL ELEMENTS : Lithium, Sodium. 

Sub-group A (similar to the typical elements) : Potassium, 

Rubidium, Caesium . 
Sub-group B : Copper, Silver, Gold. 


Group I in the Periodic Table includes Lithium and 
Sodium, with Potassium, Rubidium and Csesium in section A 
and Copper, Silver and Gold in section B. The first five of 
these elements are called the Alkali Metals. They form a 
well-marked group of metals exhibiting great similarity of 
chemical and physical nature, and this similarity extends to 
their compounds as well, as shown in the following table : 











Atomic weight 






Valency . . 






Specific gravity 0-59 





Atomic volume 









62 -5 - 



Atomicity . 






Colour of va- 









Simplest oxide 
Peroxide . . 



K,0 4 

Rb,0 t ; 
Rb,0 4 


Chloride . 






Carbonate . 
Hvdnde . 

Li 2 CO, 

Na 8 CO. 

K KH' 


Ca CO, 

Action on water 









very quickly 

very quickly 


and ignites. 

and ignites. 

and ignites. 











The alkali metals are strongly electropositive, and their 
chemical activity increases with increase of atomic weight. 
They readily oxidize in the air and decompose water in the 
cold forming hydroxides, with evolution of hydrogen 

2R + 2H 2 O = 2R.OH + H 2 . 

They are soft metals and can easily be cut with a knife or 
moulded between the fingers. Owing to their affinity for 
oxygen, they are kept in sealed tins or under rock-oil, a liquid 
that contains no oxygen. 


Group in Periodic System : I ; Symbol : Li ; Valency : 1 ; 
Atomicity of Vapour: 1 ; Atomic Weight: 6-94 ; Melting-Point : 
180 : Specific Gravity : 0-59. 

Although lithium is widely distributed in nature, it is found 
only in small quantities. Its chief ores are lepidolite and 
spodumene (double silicates of aluminium and lithium, 
generally containing sodium and potassium silicates as well). 
It is also found in the waters of springs at Redruth (Cornwall) 
and other places. 

It is a silvery-white metal which rapidly tarnishes in moist 
air. It will burn if heated in air or in oxygen, forming 
lithium monoxide, Li 2 0. This is a white solid soluble in 
water, with which it forms lithium hydroxide 
Li 2 + H 2 O = 2L10H. 

Lithium chloride, LiCl, closely resembles sodium chloride, 
but lithium carbonate, Li 2 C0 3 , is only slightly soluble in 
water and on heating splits up into lithium oxide and carbon 

Li 2 C0 3 = Li 2 + C0 2 , 

in these respects differing from the carbonates of the other 
alkali metals. 

Lithium compounds when heated in the Bunsen flame 
colour it a deep red ; this property is made use of for the 
detection of lithium. 

Since lithium urate is the most soluble urate known, a 
solution of a lithium compound is often taken as a medicine 

Nat. Portrait Gattery* 




in cases of uric acid diseases (e.g. rheumatism ), but it is of 
doubtful efficacy. 


(jfroup in Periodic System : I ; Symbol : Na ; Valency : 1 ; 
Atomicity of Vapour : 1 ; Atomic Weight : 23-0 ; Melting- 
point : 96 ; Specific Gravity : 0-97. 

Occurrence. The chief naturally occurring compounds 
of sodium are sodium chloride, which is found in the solid state 
as rock-salt in England, Germany, Galicia and many other 
countries, and in solution in brine-springs (Droitwich, etc.) 
and in the sea ; Chile saltpetre or sodium nitrate, beds of which 
occur in South America ; and a sodium carbonate or sesqui- 

carbonate found in East 

Africa, Australia, and in 
smajler quantities in 

many other hot countries . 
Qali (whence our word 
alkali) was the crude 
sodium carbonate ob- 

tained by the Arabian 

, , . / f Al 
alchemists from the ashes 

of maritime plants. 
Manufactur e. 

Sodium was first ob- 

FIG. 45.-Davy's Apparatus for the tained b ? Sl * HUMPHRY 
Preparation of Sodium. DAVY in 1807 by the 

electrolysis of fused 

caustic soda (NaOH), and, after an interval during which 
it was prepared by other means, it is still prepared by an 
elaboration of Davy's method. 

In CASTNER'S process, chiefly used in England and in 
France, the caustic soda is fused in an iron pot, through a hole 
in the bottom of which an iron rod, the cathode, rises. This 
is kept in position by allowing the caustic soda round its baso 
to solidify. < The cylindrical anode of nickel surrounds the 
upper part of the cathode, and between the two there is a 
cylinder of wire -gauze. The caustic soda is maintained in 
the fused state partly by the heat developed during electro- 



lysis and partly by a ring of gas-burners underneath the pot. 
Hydrogen and sodium are liberated from the cathode and the 
molten sodium rises into the receptacle D, from which it is 
ladled out from time to time with a perforated spoon, the 
holes of which are of such a size that the sodium is retained but 
that any caustic soda accidentally taken up at the same time 
runs through into the pot again. The hydrogen, and the oxy- 
gen which is evolved at the anode, are valuable by-products 
and are collected. 


FIG. 4G. Electrolysis of Fused Caustic Soda. 

A, Iron cathode. B. Nickel anode. C. Fused caustic soda. D. Sodium. E. Solid 
caustic soda. F. King of gas burners. 

Sodium is used in the manufacture of sodium cyanide (for 
gold-extraction), sodium peroxide, various dye-stuffs, etc. 
The annual production runs into thousands of tons. 

Properties. M.P. 96. B.P. 877. Sodium is a soft, 
silvery metal, which is readily attacked by moist air. A 
thin coating of monoxide (Na 2 0) is first formed ; this is then 
converted into hydroxide (NaOH) by the moisture present, 
and finally sodium carbonate is formed, owing to the action 
of atmospheric carbon dioxide upon the hydroxide 


2NaOH + C0 2 = Na 2 C0 3 + H 2 O. 

When burnt in the air, sodium is converted into sodium 
peroxide, Na 2 O 2 . 

Sodium readily dissolves in warm mercury, forming sodium 
amalgam, which is liquid at ordinary temperatures when the 
concentration of sodium in it is small, but forms a crystalline 
solid with higher concentrations of sodium. 

COMPOUNDS OF SODIUM. Sodium forms two oxides, Na t O 
and Na 2 2 . 

Sodium monoxide, Na 2 O, is a white crystalline solid 
formed by heating sodium peroxide with sodium or by partially 
oxidizing sodium in pure oxygen and then distilling off the 
unchanged metal. It readily combines with water, yielding 
caustic soda 

Na 2 O + H 2 O = 2NaOH. 

At 200 it reacts with hydrogen, forming sodium hydroxide 
and hydride 

Na 2 O + H 2 - NaOH + NaH. 

Sodium peroxide, Na 2 2 , is important commercially as 
an oxidizing agent. It is manufactured by heating sodium 
on aluminium trays in iron tubes in a current of dry air from 
which all carbon dioxide has been removed. It is a yellowish 
solid, readily reacting with water with evolution of oxygen 

2Na 2 2 + 2H 2 O = 4NaOH + O 2 . 

It is a powerful oxidizing agent and is used for bleaching 
and for obtaining a supply of oxygen in closed spaces such 
as submarines and diving-bells. The commercial name is 
" oxone." With cold absolute alcohol it yields sodium 
hydroperoxide or sodyl hydroxide, Na.O OH. 

Sodium Hydroxide, or Caustic Soda, NaOH. Caustic 
soda is made in three ways 

(i) By boiling a 10 per cent, solution of sodium carbonate 
with lime (Gossage's method). 

Na a CO 8 + Ca(OH) 2 ^ CaC0 3 + 2NaOH. 

The calcium carbonate is precipitated. As the action is 
reversible, the working conditions have to be carefully 


adjusted in order to get the best yield. The solution of 
caustic soda is filtered off as soon as test portions treated with 
dilute hydrochloric acid give no carbon dioxide, thus showing 
that all the sodium carbonate has been decomposed. It is 
then concentrated in specially constructed vacuum evapor- 
ators, heated by steam ; the remaining water is driven off by 
heating in polished cast-iron soda-pots. The temperature 
required to get rid of the last traces of water is above the 
melting-point of caustic soda ; the latter is therefore left in the 
molten state and is run off and either sealed up in metal drums 
or cast into the sticks or pellets so familiar in the laboratory, 
(ii) By heating a mixture of sodium carbonate and iron ore 
(Fe 2 O 3 ) in revolving cylinders, when sodium ferrite and carbon 
dioxide are formed : 

Na 2 C0 3 + Fe 2 3 = 2NaFe0 2 + CO 2 . 

The sodium ferrite is then decomposed with water and the 
precipitated ferric oxide allowed to settle : 

2NaFeO 2 + H 2 = 2NaOH + Fe 2 O 8 . 

(iii) By Electrolysis. The electrolytic processes may be 
divided into three classes 

(a) Electrolysis of sodium chloride solution, using a 

diaphragm in the cell. 
(6) Electrolysis of sodium chloride solution, in a cell 

without a diaphragm. 

(c) Electrolysis of fused sodium chloride, and treating 
the resulting lead-sodium amalgam with water (p. 199). 

In each case chlorine is obtained as an equally important 

(a) THE HARGKEAVES-BIRD METHOD. The sides of the cell 
in this process are made of compressed asbestos impregnated 
with sodium silicate, and although liquid diffusion cannot 
take place through them they are permeable to ions. The 
cell is filled with salt-solution which enters at the bottom, the 
liquid after electrolysis escaping through a pipe at the top. 
The anodes are made of carbon, or of rods of lead carrying 
lumps of gas-carbon, and dip into the solution in the cell. 
The cathode consists of a net of copper wire covering the 



outer surface of the asbestos sides, which constitute the 
diaphragm. The whole cell is enclosed in an outer chamber 
into which steam can be blown. On electrolysis, chlorine is 
liberated at the anode and escapes through a pipe at the top, 
while sodium ions pass through the asbestos diaphragm and 
are discharged at the cathode. The sodium thus formed is 
converted into a solution of caustic soda by the steam which 
is blown in. 




Salt -sokttion. 

Caustic Soda -solution 
FIG. 47. Hargreaves-Bird Cell. 

Other diaphragm processes have been invented by Le 
Sueur, etc. 

(6) THE CASTNEII-KELLNER PROCESS. In this process the 
cell consists of a rectangular tank divided into three com- 
partments by non- porous partitions which reach nearly but 
not quite to the floor. On the floor is a layer of mercury of 
such a depth that the three compartments are completely 
separated from one another by it. 

The middle compartment contains water, into which dip 
a number of iron rods, forming the cathode. The end cham- 
bers are filled with strong salt solution and contain the anodes, 


which are made of carbon rods. When the current is passed 
chlorine is evolved at the anodes and is led away and com- 
pressed ; sodium ions pass to the mercury where they are 
discharged, the sodium dissolving in the mercury. By 
means of an eccentric the cell can be gently rocked, and this 
movement causes circulation of the sodium amalgam, which 
when it enters the middle chamber reacts with the water and 
forms a solution of caustic soda, hydrogen coming off from 
the cathode. Continuity of this process obviously depends 
upon continuous circulation of the mercury. In the above 
case this is brought about by rocking the cell, but other 
methods also are used. 

FIQ. 48. Castner-Kellner Cell. 

(c) Sodium hydroxide has also been manufactured by the 
electrolysis of fused sodium chloride, molten lead being used as 
the cathode. The lead-sodium amalgam thus formed was 
decomposed by steam. The practical difficulties of this 
method were very great, largely owing to the high melting- 
point of sodium chloride. 

Properties. Caustic soda is a white crystalline solid which 
deliquesces in moist air. On prolonged standing, however, 
the liquid obtained goes solid again owing to the absorption 
of carbon dioxide from the air, resulting in the formation of 
sodium carbonate, 

2NaOH + CO, = Na 2 C0 3 + H 2 0. 

Caustic soda is very soluble in water, heat being evolved. 
The solution is strongly alkaline ; caustic soda, indeed, is one 


O rf the strongest bases known, i.e. it has a high degree of 
(ionization in aqueous solution. The solution will precipitate 
the hydroxides (sometimes the oxides) of heavy metals from 
solutions of their salts. 

In the laboratory, caustic soda is used as a reagent in 
qualitative analysis (preparation of hydroxides), for the 
absorption of carbon dioxide and sulphur dioxide from gas 
mixtures, for the titration of acids in volumetric analysis, 
and for the saponification or hydrolysis of esters and other 
organic compounds. 

The chief impurity in caustic soda is iron (from the evapor- 
ating vessels). Caustic soda may be purified by recrystalliza- 
tion from alcohol, and is then described as " pure by alcohol." 

In commerce, caustic soda is largely used for soap-making. 
Ordinary white soap consists chiefly of sodium stearate and 
palmitate, while " soft soap " consists of the corresponding 
potassium salts. Caustic soda finds also many other applica- 
tions, especially in the dye industry and in the preparation of 
artificial silk. 

Sodium hydride, NaH, is a white crystalline solid pre- 
pared by passing dry hydrogen over heated sodium. It is 
decomposed by water, yielding hydrogen and caustic soda : 

NaH + H 2 = NaOH + H 2 . 

Sodium carbonate, Na 2 CO 3 . 10H 2 0. Sodium carbonate, or 
" soda," is one of the most important compounds of sodium. In 
the form of a compound with sodium bicarbonate (NaHC0 3 ), 
called sodium sesquicarbonate, Na 2 C0 3 .NaHC0 3 .2H 2 0, it 
occurs in vast quantities in East Africa, but most of the soda 
of commerce is manufactured from common salt. The oldest 
of the processes now or formerly employed was invented by 
NICHOLAS LEBLANC (1742-1806), who first converted sodium 
chloride into sodium sulphate and then heated the latter with 
carbon and calcium carbonate. He was a doctor, and his 
wealthy patient, the Duke of Orleans, lent him 8,000 to 
work the process on a commercial scale. Then came the 
Revolution. The Duke was guillotined in 1793, and Leblano 
was ordered to make his process public, as f>oda supplies in 



France were running very short, owing to the blockade. 
Reduced to poverty, Leblanc lived on in misery till 1806, 
^hen he committed suicide. 

In England, the manufacture of soda was first established 
by MUSPBATT in 1824, who used the Leblanc process, but had 
some difficulty in persuading the public that his product 
really was soda. (The idea of the chemical individual is still 
grasped only with the greatest difficulty by the layman, who 
most often regards the product of the laboratory as an excel- 
lent imitation, but not the real stuff !) 

The Leblanc process now has very serious rivals, and is 
nearly, if not quite, obsolete. It has, however, been long in 
dying, since the hydrochloric acid obtained as a by-product 
is very valuable. It is interesting to remember that in the 
early days this acid was allowed to escape into the atmo- 
sphere, and proved such a nuisance that an Act of Parlia- 
ment was passed to compel manufacturers to absorb it. 
It would need an Act of Parliament now to prevent them 
from absorbing it. 

MODERN LEBLANO PROCESS. (i) A charge of salt is first 
mixed with an approximately equal weight of concentrated 


FIG. 49. Saltcake Furnace Old Form. 

P. Furnace. S. Pan for first stage of process. The arrows Indicate the path of the 

hot gases. 

sulphuric acid, and heated, when hydrochloric acid and sodium 
hydrogen sulphate are formed, 

NaCl + H 2 SO 4 - NaHSO, + HC1. 

The residue, consisting of the sodium hydrogen sulphate 
and unchanged salt, forms a pasty mass which is raked out 




on to a hearth heated to a much higher temperature, when 
sodium sulphate is formed with further liberation of hydro- 
chloric acid 

NaHS0 4 + NaCl - Na 2 S0 4 + HC1. 

In the older forms of furnace used for this purpose, the 
hearth was heated directly by hot furnace gases, but as these 

Fia. 50. Saltcake Furnace New Form. 

S. Pan or first stage of process. F. Furnace for heating S. F'. Furnace for heating 
hearth. 0. Exit flue for hot gases from F'. E. Exit pipe for the hydrochloric acid ga. 

mixed with the hydrochloric acid gas formed and thus ren- 
dered it impure, the hearth is nowadays arranged in such a 
way that there is a brick partition between it and the hot 

The hydrochloric acid gas is absorbed in towers packed with 

FIG. 51. Black Ash Revolving Furnace. 

A. Furnace for generating hot gases. B. Cylindrical revolving furnace. CC. Evapor- 
ating pans. DE, Cog wheels for driving B. F. Opening to furnace. 

coke over which a stream of water trickles, a solution of 
hydrochloric acid running out at the bottom. These towers 
were invented by WILLIAM GOSSAGE in 1836. 

(ii) The sodium sulphate, or salt-cake, is then mixed with 
powdered limestone (CaCO) and small coal, and heated 



strongly in cylindrical iron furnaces which revolve around 
their longitudinal axes. Each furnace will hold about 18 
tons, and is heated by hot gases which enter it through an 
opening at the end. After passage through the furnace, the 
hot gases are led over the surface of soda solutions in evaporat- 
ing pans and their heat thus usefully employed in evaporation. 

During the operation, the furnace is slowly rotated, thorough 
mixing and heating being thus ensured. The end of the 
reaction is marked by the appearance of flames of carbon 
monoxide, and the semi-liquid mass is then run out through 
an opening in the furnace into iron trucks below, where it 

The product is called black-ash, and consists of sodium 
carbonate mixed with excess of limestone and coal, sodium 
and calcium sulphides, lime, and other substances. The 
reaction in the furnace may be represented by the following 

(i) Na 2 SO 4 + 40 - Na 2 S + 4CO 
(ii) Na 2 S + CaC0 3 = CaS + Na 2 C0 3 . 

The black- ash is crushed and lixiviated or extracted with 
water, in the cold, in a series of iron tanks which are arranged 
in such a way that the maximum concentration of solution is 
obtained, together with complete extraction. The principle 
of this method will be clear from Fig. 52. 

Direction of movement of solution. 

FIG. 52. Freshness of black-ash. 

A-E represent five tanks in which black-ash is placed. 
Tank A is now filled with water, when sodium carbonate 
solution is formed, but the solution is not very concentrated. 


Tank B is then filled with the soda solution from A, which 
dissolves more soda. Meanwhile, A is filled with water again, 
and when tank C is filled with the soda-solution from B, B is 
filled with the solution from A, and so on. This process is 
continued until the black-ash in A is completely exhausted 
of sodium carbonate, when fresh black-ash is put in and the 
exhausted ash taken out. By this time, the original solution 
from A has arrived at E and is probably saturated with soda. 
It is therefore run off for evaporation ; but if not, it is returned 
to A, which is now charged with the fresh black-ash. Here 
it will become saturated and may therefore be run off. 

Thus the principle used is that of adding fresh water to 
black-ash which is nearly exhausted (in this way completing 
the extraction of the sodium carbonate), while the fresher ash 
is extracted by solutions, the fresher the ash the greater the 
concentration of the solution used to extract it (in this way 
obtaining a maximum concentration of the solution). This 
principle is of very general use in manufacturing chemistry, 
and is called the Principle of Counter-Currents, since, in this 
case, for example, the strength of the solution in sodium 
carbonate increases as we pass from left to right in the 
diagram, while the richness of the black-ash in sodium 
carbonate is greater as we pass from right to left. 

The solution of sodium carbonate is allowed to stand, in 
order that insoluble impurities may settle, and is then decanted 
off and evaporated in shallow pans, when crystals of sodium 
carbonate decahydrate, or washing soda, Na 2 CO,.10H 2 0, 

The residue left in the extraction tanks is called alkali 
waste, and is commercially valuable on account of the sulphur 
it contains. It is only by making use of the by-products that 
the Leblanc process is able successfully to compete with the 
other main process of manufacturing sodium carbonate ; 
hence the sulphur is extracted from the alkali-waste by the 
CHANCE process (see p. 457), and the hydrochloric acid 
obtained in the salt-cake stage carefully absorbed, as already 



was first used in 1836 by THOM, but is sometimes called the 
SOLVAY process since it was revived with many improvements 
by this chemist in 1865. The reaction involved is delightfully 
simple ; here, as in many other cases, it was the practical 
difficulties of carrying out the principle on a large scale that 
required much labour and genius to overcome them. 

When ammonium bicarbonate is added to a concentrated 
solution of sodium chloride, sodium bicarbonate is precipi- 

NH 4 HC0 8 + NaCl ^= NaHC0 3 * + NH 4 C1. 
If sodium bicarbonate is heated, sodium carbonate is formed 

2NaHCO 3 - Na 2 CO 3 + CO 2 + H 2 O. 

In practice, a concentrated solution of salt is taken and 
saturated with ammonia, when any impurities that are 
present separate and are allowed to settle. The ammoniacal 
salt-solution is then pumped into towers (" carbonators ") 
where a stream of carbon dioxide is blown through it. The 
following reactions then occur 

(i) 2NH 3 + CO a + H 2 O ;= (NH 4 ) 2 C0 8 

(ii) (NH 4 ) 2 C0 3 + H 2 + C0 2 ;== 2NH 4 HC0 3 

(iii) NH 4 HC0 3 + NaCl ^= NaHC0 3 + Nfc 4 Cl. 

The conditions of these reversible reactions are adjusted so 
as to give the maximum yield of sodium bicarbonate. Heat 
is evolved in the reactions and the carbonators are therefore 
cooled by currents of water on the outside. 

The liquid which flows out at the bottom of the towers 
carries the precipitated sodium bicarbonate with it, and is run 
over rotating vacuum filters where the sodium bicarbonate is 
retained. This is scraped off and heated in special furnaces, 
which differ in construction in different factories, and anhy- 
drous sodium carbonate, or soda-ash, is left. 

2NaHC0 3 = Na 2 C0 3 + H 2 + C0 2 . 

One of the chief points of interest in this process is its self- 
contained nature. Thus, the carbon dioxide is prepared by 
heating lime-stone in a lime kiln 

CaC0 8 - CaO + C0 a , 


and the other product, lime, is used to regenerate the ammonia 
from the ammonium chloride formed in the reaction given in 
equation (iii) above 

CaO + 2NH 4 C1 = CaCl 2 + 2NH 8 + H 2 O. 

The carbon dioxide evolved when the sodium bicarbonate 
is heated is of course used again ; the equations show that in 
a single operation, half the original amount of carbon dioxide 
is converted into sodium carbonate and the other half regained 
as gas. 

The only waste product is therefore calcium chloride, and 
in spite of desperate efforts to persuade local authorities to 
use this to keep their roads damp and free from dust in warm 
weather, it remains a drug on the market. However, the 
soda industry is fairly profitable in spite of this intractable 
substance, for the leading British soda-manufacturers paid a 
dividend of 100 per cent, for many years. 

Sodium carbonate is also manufactured to some extent 
electrolytically, or rather by the action of carbon dioxide 
upon solutions of caustic soda obtained electrolytically. 

Sodium carbonate is used in the manufacture of glass, soap, 
and paper, and for softening water. Washing-powders con- 
sist mainly of a mixture of sodium carbonate and powdered 
soap. Bath salts are dyed and scented soda crystals. 

Properties. Anhydrous sodium carbonate is a white solid 
which readily takes up water to form crystalline hydrates. 
The decahydrate, Na 2 C0 3 .10H 2 O, is washing-soda. In dry 
weather, washing-soda crystals effloresce ; the resulting white 
powder is the monohydrate, Na 2 CO 3 .H 2 0. In the labora- 
tory, sodium carbonate is used as a mild alkali, since when 
dissolved in water it is partially hydrolysed 

(i) Na 2 C0 3 ^z Na + Na + C0 8 " 
(ii) H 2 ;= H' + OH' 
(iii) C0 3 " + 2H' + 20H' ^= H 2 CO 8 + 20H'. 

It first ionizes into sodium ions and C0 3 " ions, but the latter 
immediately combine with the few hydrogen ions normally 
present in water, to form un-ionized carbonic acid ; more water 


then ionizes to restore the equilibrium H* + OH' v N H 2 0, 
but as soon as hydrogen ions are formed they react with the 
CO 3 " ions to give more carbonic acid, and this process goes on 
until equilibrium has been set up between the various sub- 
stances present. When this occurs, there will be a number of 
hydroxyl ions (OH') in excess equal to the number of hydrogen 
ions that have been seized by the C0 3 " ions to form carbonic 
acid, therefore the solution will have an alkaline reaction. 

The strength of a solution of sodium carbonate may be 
estimated by titration with standard acid, if the solution be 
kept boiling so that the carbon dioxide produced is driven 
off, or if an indicator insensitive to carbonic acid (e.g. methyl 
orange) be used. Of these alternatives, the former is prefer- 

Sodium bicarbonate, NaHC0 3 , is a white crystalline 
solid with a very slightly alkaline (practically neutral) reaction 
when dissolved in water. It is used as a baking-powder, 
dnce on heating it loses carbon dioxide, which makes the 
dough " rise." It is also used in making " sherbet " powder. 
Sodium chloride, NaCl, or common salt, is a white solid 
crystallizing in anhydrous cubes. The occurrence of salt 
is mentioned later, on p. 497. Salt is made from brine 
by evaporation, usually by steam-heat and under reduced 
pressure. By arranging the evaporation chambers in a series 
in such a way that the steam from the boiler passes first into 
the chamber in which the pressure is highest and thence into 
chambers in which the pressure is lower, the maximum 
efficiency is obtained. This arrangement is called the Multiple 
Effect Vacuum Evaporation Plant. 

In hot countries near the sea, salt is obtained by the evapora- 
tion of sea-water (e.g. in Spain and the South of France). 
In cold countries, on the other hand, the sea- water is first 
frozen, and the residual concentrated brine then evaporated. 
Pure sodium chloride may be obtained by passing hydro- 
chloric acid gas into a saturated solution of the impure sal^ 
when the great increase in the active mass of the chlorine ions 
causes the solubility-product (p. 135) of the salt to be exceeded, 
while that of the more soluble impurities (chiefly magnesium 



chloride) is not reached. Salt is therefore precipitated while 
the impurities remain in solution. 

The uses of salt are well known. It forms the starting- 
point for the production of practically all the sodium and 
chlorine and their compounds used in commerce. 

Sodium bromide, NaBr, and sodium iodide, Nal, 
resemble the corresponding potassium salts (pp. 525, 213), and 
may be made in similar ways. 

Sodium hypochlorite, NaOCl, is described under hypo- 
chlorous acid (page 515). 

Sodium sulphate, Na 2 S0 4 , is obtained commercially as 
salt-cake in the Leblanc sodium carbonate process (p. 202). 
When an aqueous solution of sodium sulphate is evaporated 
the salt separates as the decahydrate, Na 2 S0 4 .10H 2 O ; this 
is known as GLAUBER'S salt and is used as a purgative. The 
crystals of the decahydrate gradually effloresce (p. 444) in 





Temperature . 

4-0 50 60 

FIG. 53. 


8O 90' 100' 

dry air. The solubility of Glauber's salt in water gradually 
increases up to a temperature of 32-5 ; at this temperature, 
however, the solid decahydrate spilts up into water and 


anhydrous sodium sulphate which is peculiar in becoming 
less soluble as the temperature rises. The solubility curve 
of sodium sulphate therefore exhibits a sharp break at 32-5. 

It should be carefully noted that this curve is really com- 
posed of two distinct curves, (i) the solubility curve of the 
decahydrate from to 32-5, and (ii) that of the anhydrous 
salt above 32-5. There is no sharp break, of course, in the 
constitution of the solution at 32-5 ; the only change is in 
the nature of the solid with which the solution is in contact. 1 

Sodium bisulphate, NaHSO 4 , is a white crystal-line solid 
with an acid reaction in aqueous solution. It is obtained 
commercially as a by-product in the manufacture of nitric 
acid from caliche (see below), and is then known as nitre-cake. 
When heated it loses water and is converted into sodium 

2NaHS0 4 = Na 2 S 2 O 7 + H 2 O. 

Sodium thiosulphate, Na 2 S 2 3 .5H 2 0. See p. 478. 

Sodium nitrate, NaN0 8 , occurs in large quantities in 
South America and is commonly known as Chile saltpetre ; 
the crude material is known locally as caliche. The sodium 
nitrate is purified by solution in water and recrystallization ; 
the mother liquors contain sodium iodate and are used as a 
source of iodine (p. 529). 

Sodium nitrate is a white deliquescent crystalline solid. 
When strongly heated it loses oxygen and is converted into 
sodium nitrite, NaNO 2 

2NaN0 3 = 2NaNO a + 2 . 

Sodium nitrate is very largely used (some millions of tons 
per annum] as a fertilizer for corn lands. It is applied to the 
growing crop as a top-dressing in the spring ; its effects last 
for one year only, as it is so soluble that it is all leached out 
of the soil by the winter rains. It is also used in the manu- 
facture of nitric acid. 

1 The left-hand curve represents the equilibrium between the decahy- 
drate and the saturated solution ; the right-hand curve represents the 
equilibrium between the anhydrous salt and the same saturated solution. 


Sodium nitrite, NaNO 2 , is chiefly made at the present 
day by the action of oxides of nitrogen (obtained synthetically 
from the air, p. 354) upon caustic soda (p. 355). It is also 
made by fusing sodium nitrate with lead 

NaN0 3 + Pb = NaN0 2 + PbO, 

and by passing a current of sulphur dioxide through a hot 
concentrated solution of sodium nitrate mixed with lime 

NaN0 3 + CaO + S0 2 = NaN0 2 + CaS0 4 . 

It is a pale yellow crystalline solid, somewhat deliquescent 
and readily soluble in water. It is chiefly employed in the 
manufacture of aniline dyes. For the properties of nitrous 
acid and the nitrites see p. 367. 

Sodium hydrogen sulphide, NaHS, may be made by 
passing sulphuretted hydrogen into caustic soda solution until 
no more will dissolve 

NaOH + H 2 S = NaHS + H 2 O. 

If to the solution of the sodium hydrogen sulphide an 
equivalent proportion of caustic soda be added, the normal 
sulphide, Na 2 S, is formed 

NaHS + NaOH = Na 2 S + H 2 0, 

and may be obtained in the form of white hydrated crystals, 
Na 2 S.9H 2 O, by evaporating the liquid to crystallization. 

It is made commercially in the anhydrous state by strongly 
heating sodium sulphate with powdered coke 

Na 2 S0 4 + 2C = Na 2 S + 2C0 2 . 
Polysulphides of sodium, Na 2 S a; , are known. 
Other sodium salts have been described in various places 
in this book. See the index. 


Group in Periodic System : I ; Symbol : K ; Valency : 1 ; 
Atomicity of Vapour : 1 ; Atomic Weight : 39-10 ; Melting 
Point : 62-3 ; Specific Gravity : 0-865. 

History. The Arab chemists gave the name alkali (i.e. 
" the ash ") to the ashes obtained by burning plants. Land 


plants give an ash containing much potassium carbonate, 
while the ash of maritime plants contains more of the sodium 
salt than of the potassium. The distinction in properties 
between these two alkalis was not clearly observed until the 
seventeenth century, when the difference between potash and 
soda was noticed by BOHN. It was later discovered that soda 
can be made from common salt, and hence it was called 
mineral alkali, while potash was called vegetable alkali. 
Vegetable alkali was purified by extracting wood-ashes with 
water, allowing insoluble matter to settle, and evaporating 
the clear solution in pots ; hence the name potash. The 
somewhat aristocratic name potassium is therefore of very 
plebeian origin. 

The preparation of caustic alkalis from the " mild " vege- 
table and mineral alkalis by treatment with lime was known 
in the thirteenth century, and caustic soda and potash were 
considered to be elements by D ALTON. LAVOISIER, however, 
with his usual wonderful foresight, regarded them as oxygen 
compounds a forecast that was brilliantly vindicated by 
the isolation of metallic sodium and potassium by DAVY in 

Occurrence. Potassium compounds are widely distri- 
buted in nature. Igneous rocks generally contain potassium 
silicate in the form orthoclase or potash felspar, KAlSi 3 8 , or 
as leucite, KAlSi 2 6 , or muscovite, KH 2 Al 3 (Si0 4 ) 8 . From 
these minerals, potassium compounds find their way into 
clays and soil, where they form one of the essential substances 
for the growth of plants. 

The principal source of the potassium salts of commerce is 
the vast bed of saline deposits at Stassfurt, in Prussia, although 
the deposits in Alsace are likely to prove of nearly equal im- 
portance. The Stassfurt deposits are considered to consist of 
the salts left by evaporation of an inland sea. The main 
potassium compound in them is carnallite, KCl.MgCl 2 .6H 2 0, 
from which most of the world's supply of potassium and its 
compounds is made. 

Preparation. Potassium is of less commercial importance 
than sodium, and is therefore not extracted on a very large 


scale. The process employed for the preparation of the metal is 
similar to that used in the manufacture of sodium, viz., 
electrolysis of the fused hydroxide. 

Properties. In general chemical and physical properties 
potassium closely resembles sodium ; it is, however, distinctly 
more reactive. Potassium has a specific gravity 0-865, 
melting-point 62-3, and boiling-point 757. The vapour is 
brilliant green in colour and consists of monatomic molecules. 


Potassium hydride, KH, is made by passing hydrogen 
over heated potassium. It is a white crystalline solid, which 
reacts with water to give caustic potash and hydrogen 

KH + H 2 = KOH + H 2 , 

and with carbon dioxide to form potassium formate 
KH + C0 2 = H.COOK. 

Potassium monoxide, K 2 0, can be made by heating 
potassium nitrate with potassium. It dissolves in water to 
form potassium hydroxide, and in dilute acids to form 
potassium salts. 

Potassium tetroxide, K 2 O 4 , is formed as a yellow powder 
when potassium burns in air or oxygen. It reacts vigorously 
with water, yielding caustic potash and oxygen, and, in the 
cold, hydrogen peroxide 

K 2 4 + 2H 2 = 2KOH + H 2 2 + 2 . 

Potassium hydroxide, KOH, is very similar to caustic 
soda, and is made in similar ways. 

Potassium carbonate, K 2 C0 3 , is manufactured from 
carnallite in the following way. The carnallite is first melted 
and then allowed to cool, when potassium chloride crystallizes 
out and is removed. A concentrated solution of the potassium 
chloride is then made and this is mixed with hydrated mag- 
nesium carbonate. On passing a stream of carbon dioxide 
through the liquid a precipitate of a double salt 

KHCO,.MgCO,.4H 2 0, 
is obtained 


3(MgC0 3 .3H s O) + 2KC1 + CO, 

= 2(KHC0 3 .MgC0 3 .4H i! 0) + MgCl 2 . 

This solid is removed and treated with magnesium oxide, 
when the hydrated magnesium carbonate is re-formed and 
potassium carbonate left in solution. After filtration, the 
solution of potassium carbonate is evaporated to crystallization. 

2(KHC0 3 .MgC0 3 .4H 2 0) + MgO 

- 3(MgC0 8 .3H 2 0) + K 2 C0 8 . 

" Commercially pure " potassium carbonate is commonly 
known as pearl-ash. 

Potassium carbonate is also manufactured from the 
potassium chloride (obtained from carnallite) by the Leblanc 
process (p. 201). The Solvay process (p. 204) cannot be used 
on account of the high solubility of potassium bicarbonate. 

Potassium carbonate is a white deliquescent crystalline solid, 
very soluble in water, from which it crystallizes as the 
dihydrate, K 2 C0 3 .2H 2 0. 

Potassium bicarbonate, KHC0 3 , is made by passing 
carbon dioxide into a solution of the normal carbonate 

K 2 C0 3 + H 2 + C0 2 = 2KHC0 3 . 

It is much more soluble than the corresponding sodium 
salt, which otherwise it very closely resembles. 

Potassium chloride, KC1, occurs at Stassfurt and in 
Alsace as sylvine, KC1, and at Stassfurt also as carnallite. It 
is made from carnallite as described under potassium car- 
bonate, or by treating the carnallite with the hot mother- 
liquors from other recrystallizations. On cooling, the potas- 
sium chloride separates and is purified by recrystallization. 

Potassium bromide, KBr, is a white crystalline solid, used 
in photography and medicine. For its preparation see p. 525. 

Potassium iodide, KI,is very similar to the bromide. 
Its aqueous solution is a good solvent for iodine (p. 530). 

Potassium chlorate, KC10 3 , is manufactured by the 
electrolysis of a hot concentrated solution of potassium 


chloride. 1 The caustic potash formed at the cathode reacts 
with the chlorine liberated at the anode, according to the 

6KOH + 3C1 2 = 5KC1 + KC10 8 + 3H 2 0. 

Potassium chlorate is a white anhydrous crystalline solid. 
It is a strong oxidizing agent, and is commonly used in the 
laboratory as a source of oxygen. Its oxidizing powers render 
it a good germicide ; it is therefore sometimes made up into 
" chlorate lozenges " for the cure of sore throat. Mixtures 
of potassium chlorate and sulphur or red phosphorus explode 
very violently when rubbed or struck. 

Potassium sulphate, K 2 S0 4 , forms anhydrous crystals 
(cf. sodium sulphate). It is largely used as a fertilizer in 
agriculture, especially for corn, of which it greatly improves 
the straw, and for potatoes. For the double salt, alum y of 
potassium sulphate and aluminium sulphate, see p. 280. 

Potassium nitrate, KN0 3 , or nitre, is found in the soil of 
certain tropical countries, such as India (" Indian saltpetre "). 
It is made by the bacterial oxidation of nitrogenous animal 
refuse in the presence of wood-ashes ; this process is a regular 
industry in India. 

It is also manufactured from Chile saltpetre, NaN0 3) by 
adding this compound to a hot concentrated solution of 
potassium chloride, when " double decomposition " occurs 

NaN0 3 + KC1 = KN0 3 + NaCl. 

By regulating the concentrations, the comparatively 
sparingly-soluble sodium chloride can be made to separate 
practically completely. After removal of the salt the solution 
is allowed to cool, when the potassium nitrate crystallizes out. 

Potassium nitrate forms white anhydrous crystals which 
are soluble in water but are not deliquescent and therefore 
can be used in the manufacture of gunpowder ; sodium 
nitrate, on the other hand, is very deliquescent and is, 
naturally, useless for this purpose. 

1 See also p. 505. 


When potassium nitrate is strongly heated it loses oxygen 
and is converted into potassium nitrite 
2KN0 3 - 2KN0 2 + 2 . 

Potassium cyanide, KCN, is manufactured by passing 
ammonia into molten potassium carbonate containing 
powdered carbon 

C + K 2 C0 3 + 2NH 3 - 2KCN + 3H 2 0. 

It is also prepared, mixed with sodium cyanide, by fusing 
potassium ferrocyanide with metallic sodium 

K 4 Fe(CN) 6 + 2Na = 4KCN + 2NaCN + Fe. 

It is a white crystalline solid, soluble in water and smelling 
strongly in moist air of prussic or hydrocyanic acid, HON. 
Its aqueous solution is alkaline to litmus, since hydrocyanic 
acid is a weak acid and hydrolysis of the salt therefore occurs. 

Potassium cyanide is used in electro-plating and in the 
chemical industry. In gold-extraction (p. 236) it has now 
been replaced by sodium cyanide, which is much cheaper 
and, bulk for bulk, contains more of the ON radical the 
essential part of the molecule. It is extremely poisonous. 

Potassium ferrocyanide, K 4 Fe(CN), or yellow prussiate oj 
potash, is made by heating a mixture of nitrogenous organic 
refuse (such as horns, waste hides and hoofs), scrap iron, and 
potassium carbonate. Blood, iron and potassium carbonate 
may be used, whence (so it is said !) the name prussiate of 
potash. The fused mass is treated with water and after 
insoluble matter has been allowed to settle the clear solution 
is run off and evaporated, when lemon-yellow crystals of 
K 4 Fe(CN) 6 .3H 2 are formed. 

When chlorine is passed into a solution of potassium 
ferrocyanide the latter is oxidized to potassium ferricyanide,, 
K 8 Fe(CN) 

2K 4 Fe(CN) 6 + C1 2 = 2KC1 + 2K 3 Fe(CN) 6 . 

Potassium ferricyanide is a dark red crystalline solid, 
sometimes known as red prussiate of potash. It and the ferro- 
cyanide are used in qualitative analysis as tests for iron in 



Ferrous Ions. 

Ferric Ions. 

Potassium ferrocyan- 

White ppt. rapidly 
going blue. 

Deep blue ppt. 
" Prussian Blue." 

Potassium ferricyan- 

Deep blue ppt. 

No ppt. Brown 

Potassium thiocyanate, KCNS, is made by heating a 
mixture of potassium cyanide and sulphur 

KCN + S = KCNS. 

It is a white, deliquescent, crystalline solid readily soluble 
in water. Its solution with a solution containing ferric ions 
gives a deep blood-red colour, usually said to be due to the 
formation of ferric thiocyanate but probably caused by the 
formation of a complex anion. The equation corresponds 
with the former explanation 

FeCl 3 + 3KCNS ;= 3KC1 + Fe(CNS),. (See p. 121.) 


1. Show that in the physical and chemical properties of the elements 
and their compounds the alkali-metals form a natural group. 

2. Describe the Leblanc process for the manufacture of sodium 

3. Describe the Solvay or Ammonia-Soda process for the manu- 
facture of sodium carbonate. Why cannot potassium carbonate be 
made in a similar way ? 

4. Give an account of the manufacture of potassium carbonate and 
potassium chlorate. 

5. Why is a solution of potassium cyanide alkaline to litmus ? How 
is this salt made, and what are its uses ? 

6. How is sodium prepared commercially ? 

7. What do you know of the history of sodium and potassium T 


GROUP I, Sub-group B 


Practically the only point of resemblance between these 
three metals and the alkali metals is the fact that they can 
behave as univalent elements. They show a considerable 
similarity between themselves, however, and also to the 
transition elements of Group VIII. It is interesting to note 
that copper, silver and gold, which form a well-defined group 
in the periodic system, are associated also in daily life from 
their use in coinage and in electroplating. 

Silver is uniformly univalent ; copper can be both univalent 
and bivalent, while gold is sometimes univalent and sometimes 

They are all very resistant to the action of water, and thus 
differ as widely as possible, in this respect, from their associates 
in the group the alkali metals. 


Group in Periodic System : I ; Symbol : Cu ; Valency f 
1 and 2 ; Atomic Weight : 63-57 ; Melting Point : 1,083 ; 
Specific Gravity : 8-95. 

History. Metallic copper occurs in many parts of the 
ancient world and has therefore been known to man from 
immemorial antiquity. The different periods in the develop- 
ment of civilization are roughly described as the Early Stone 
Age (Palceolithic), the Later Stone Age (Neolithic), the Bronze 
Age, and the Iron Age, according to the material used by man 



for making his weapons. In the Bronze Age copper and its 
alloy with tin, bronze, had superseded the flint *of the Stone 
Age, but with advance in chemical knowledge they were 
themselves partially displaced by the more difficultly worked 
iron and steel. 

It seems probable that in the early days of the Bronze Age 
copper itself was chiefly employed ; tin was comparatively 
rare in the Mediterranean basin and most of it had to be 
obtained by perilous journeys to Britain, to Persia or to China. 
However, when tin was discovered in quantity the use of 
bronze became much more general, and finally pure copper 
was scarcely ever used except by the chemists for their 
experimental work. 

The copper used by the Egyptians was mined in the penin- 
sula of Sinai, and was called khomt ; the Greeks and Romans 
obtained it chiefly from the island of Cyprus, whence the name 
aes cyprium or Brass of Cyprus, modified later into simple 
cuprum. Cyprus was a favourite resort of the goddess Venus, 
and for this and other reasons copper was often called by the 
name of the goddess and represented by her sign, $. The 
Latin GEBEB (thirteenth century) describes the metal as 
follows : 

" Venus or Copper is a Metallick Body, livid, partaking of 
a dusky Rednefs ignible (or Jujtaining Ignition) fujible, 
extenjible under the Hammer, but refujing the Cupel. . . . 
It agrees very well with Tutia [i.e. it forms an alloy (brass) 
with zinc], which citrinizeth (or Colours) it with good Yellow- 
nejs. ... It receives Infection from {harp and acute things ; 
and to eradicate that, is not an eajie, but a profound Art " 
[i.e. it is readily attacked by acids, and this property is 
shown by it even in most of its alloys]. 

Occurrence. Enormous quantities of native copper are 
found near Lake Superior, and smaller quantities in Scan- 
dinavia, Russia, Australia, China, and other places. The chief 
ores are chalcopyrite or copper pyrites, CuFeS 2 ; copper glance 
or chalcocite, Cu 2 S ; cuprite or " red copper ore," Cu 2 ; 
malachite, Cu(OH), . CuC0 3 ; and azurite, Cu(OH) 2 . 2CuC0 8 , 

In minute traces, copper is widely distributed in soil. 


Extraction. Native copper, when occurring in large 
boulders, proves rather troublesome to deal with since it is so 
tough. In cases of difficulty, huge electrolytic cells have 
been built round the boulders and filled with acidified copper 
sulphate solution ; a thin sheet of pure copper is then made 
the cathode and the boulder is made the anode. On electro- 
lysis copper passes from the anode to the cathode, where it 
is deposited. In this way the native copper is extracted and 
refined at the same time. Smaller masses are melted down 
with a little carbon and a flux to remove earthy impurities, 
and the metal then refined by electrolysis. 

Oxide and carbonate ores are mixed with a flux and 
powdered coke and reduced in a reverberatory furnace 
Cu 2 + C = 2Cu + CO. 

Most copper ores, however, contain sulphur, and the 
extraction of the copper is a difficult matter, for the details 
of which larger books must be consulted. 

The processes employed differ in different localities, but the 
main principles of most of them are as follow 

(i) The ore is first concentrated by crushing and levigation 
in a stream of water, when the ore particles are carried on by 
the water while the earthy matter is left. 

(ii) The concentrated ore is then roasted in furnaces shaped 
like the pyrites burners of the lead chamber process. During 
the roasting the iron is converted into iron oxide while the 
copper is left as cuprous sulphide. 

4CuFeS 2 + 90, = 2Cu 2 S + 6SO a + 2Fe O t . 

(iii) The roasted ore is then mixed with sand, quartz or 
silicates and heated to fusion in a reverberatory furnace, 
when the silica and iron oxide form a fusible slag of an iron 
silicate, and the cuprous sulphide, " matte " or " coarse- 
metal" is left as a molten mass below the slag. The slag is 
run off and the matte poured into water, to granulate it. 

(iv) Repetition of the above process results in a purer 
matte, called "fine-metal." This is heated on the hearth of 
a reverberatory furnace (Fig. 54) with free access of air, or in a 
Bessemer converter (p. 548) through which hot air can be 



blown. Part of the cuprous sulphide is converted into cuprous 
oxide, which then reacts with the remainder of the sulphide to 
form copper and sulphur dioxide 

2Cu 2 S + 3O 2 = 2Cu 2 + 2SO 9 

2Cu a O + Cu 2 S = 6Cu + S0 2 . 

As the mass of copper cools it gives up the sulphur dioxide 
which was dissolved (or contained) in it ; the formation of 
these bubbles of gas in the copper makes the latter appear to 

Fio. 54. Reverberatory Furnace (Short Variety). 

be covered with blisters and it is therefore called blister- 

(v) The blister-copper still contains small quantities (up 
to 5 per cent.) of sulphur, iron, etc., as impurities. It is 
melted in a furnace and after removal of the thin layer of 
slag which rises to the surface, is stirred up with poles of green 
wood. This reduces any cuprous oxide in the metal to 
copper ; the reduction is completed by addition of a little 
powdered charcoal. 

(vi) The crude copper obtained in this way is refined 
electrolytically by making it the anode in an electrolytic ceD 


containing acidified copper sulphate solution and a cathode 
consisting of a thin sheet of pure copper. 

Cathode. Pure copper. Anode. 

Pure copper. -4- Crude copper. -f- 

Gold, silver. 

Some impurities go" 4 "" Certain impurities fall to the 

into solution. bottom of the cell as anode sludge. 

The copper is transferred by the current from the anode to 
the cathode while the impurities collect on the bottom of the 
cell as a mud (gold, silver) or pass into solution and are not 
deposited (iron, nickel, zinc). The anode mud is cupelled 
(p. 228) to extract the gold and silver. 

A considerable quantity of copper is extracted from burnt 
pyrites by (a) " dry chlorination " or (6) " wet chlorination " ; 
the latter is also used to extract copper from ores that are 
too poor in copper to be worked by the main process described 

(a) Dry chlorination. The burnt pyrites are mixed with 
about 15 per cent, of common salt and heated in furnaces to a 
dull red heat. This converts the copper into cupric chloride, 
CuCl 2 , which is dissolved out in water. Iron turnings or 
pieces of scrap iron are then added to the copper chloride 
solution, when the copper is precipitated 

CuCl 2 + Fe = Cu + FeCl 2 . 

(6) Wet chlorination. The burnt pyrites or powdered ore 
is treated with a warm aqueous solution of salt and ferrous 
sulphate ; the copper passes into solution as cupric chloride, 
from which it is regained by precipitation with iron as in (a). 

Properties. Copper is a metal of a characteristic rosy-pink 
colour, which is best seen by heating the metal and reducing 
it in methyl alcohol. It is malleable and ductile at ordinary 
temperatures, although it becomes very brittle if heated nearly 
to its melting-point. It is an extremely good conductor of 
electricity and therefore is widely used in the electrical 
industry. It has a specific gravity of 8-95 ; it melts at 1,083 


and boils at 2,310. The vapour is green in colour and prob* 
ably monatomic. 

In dry air copper is stable at ordinary temperatures, but 
in moist air it is slowly converted into verdigris, a basic 
carbonate. If heated in air or oxygen it is gradually oxidized 
to black cupric oxide, CuO ; it will not, however, burn with a 
flame as probably noticed by GEBER or his translator (p. 218), 
who first wrote ignible but modified this to sustaining ignition. 

Nitric acid readily dissolves copper, with formation of 
oxides of nitrogen, e.g. 

3Cu + 8HN0 3 - 3Cu(N0 3 ) 2 + 4H 2 + 2NO, 
or Cu + 4HN0 3 - Cu(N0 3 ) 2 + 2H 2 + 2NO,. 

Concentrated sulphuric acid dissolves copper on heating, 
yielding sulphur dioxide and a mixture of copper sulphate 
and sulphides, but the dilute acid has no appreciable action. 

In the presence of air, concentrated hydrochloric acid slowly 
dissolves copper, forming cupric chloride. 

Colloidal copper can be prepared by BREDIG'S method, and 
in other ways. 

COMPOUNDS OF COPPER. Copper forms two oxides, cuprous 
oxide, Cu 2 0, and cupric oxide, CuO ; corresponding to the 
oxides are the cuprous and cupric salts in which the copper is 
univalent and bivalent respectively. 

Cuprous oxide, Cu 2 0, which occurs naturally as cuprite, 
is most conveniently prepared by reducing FEHLINU'S solution 
with grape-sugar. Fehling's solution is made by mixing 
cupric sulphate and potassium sodium tartrate in solution 
and adding caustic soda. When heated with grape-sugar it 
yields a red precipitate of cuprous oxide. This reaction is 
often used to estimate the strength of a solution of grape- 

Cuprous oxide is a red powder insoluble in water, but 
soluble in ammonia, with which it forms a complex cuprous- 
ammonium compound. It dissolves in hydrochloric acid to 
give a colourless solution of cuprous chloride, which rapidly 
turns blue in the air owing to oxidation 

Cu 2 + 2HC1 = Cu 2 Cl a + H t O. 


With sulphuric and nitric acid it yields solutions of cupric 

Cu 2 + H 2 S0 4 = CuS0 4 + H 2 + Cu. 
Cu 2 O + 2HNO 3 = Cu(N0 3 ) 2 + H 2 + Cu. 

In the latter case the copper formed is immediately dissolved 
by the excess of acid, with evolution of nitric oxide. 

Cuprous chloride, Cu 2 Cl 2 , may be made by passing 
hydrochloric acid or chlorine over copper heated in a hard 
glass tube ; we should expect to get cupric chloride in the 
latter case, but this compound is unstable, splitting up into 
cuprous chloride and chlorine at a comparatively low 

Cuprous chloride is generally made by heating a solution 
of cupric chloride with copper turnings and hydrochloric 

CuCl 2 + Cu = Cu 2 Cl 2 . 

The reaction is a reduction of the cupric chloride, and may be 
effected by many 01 her reducing agents, such as sulphur 

2CuCl 2 + 2H 2 + S0 2 = Cu 2 Cl, + H 2 S0 4 + 2HC1. 
On pouring the clear solution obtained into an excess of air- 
free water the cuprous chloride comes down as a white pre- 
cipitate, rapidly turning green on exposure to air owing to 

Cu 2 Cl 2 + oxygen + H 2 - Cu(OH) 2 .CuCl 2 . 

Cupric oxy chloride, or basic 
cupric chloride. 

Cuprous chloride is a white solid, melting at 425 and 
boiling at about 1,100 ; the density of its vapour corresponds 
to the double formula Cu 2 Cl 2 and not to the simple one CuCl 
(cf . mercurous chloride, p. 263). It dissolves in concentrated 
hydrochloric acid and also in strong ammonia solution, 
forming colourless liquids that very rapidly oxidize in the 
air and turn blue. The solution in ammonia is used as a 
solvent for carbon monoxide, with which it forms the colourless 
compound Cu 2 Cl 2 .C0.2H 2 0. When acetylene is passed 
through " ainmoniacal cuprous chloride " it gives a chocolate- 


brown precipitate of cuprous acetylide, Cu 2 C 2 , which explodes 
when dried and heated. The explosion is only slight not at 
all comparable to the violent detonation of the corresponding 
silver acetylide (p. 298) though it may be serious enough if 
large quantities explode, and fatal accidents have been caused 
by it. 

Cuprous cyanide, CuCN or Cu 2 (CN) 2 , is obtained as a 
white precipitate by adding potassium cyanide to a solution of 
cupric sulphate ; the first-formed cupric cyanide is unstable 
and loses cyanogen, leaving the cuprous salt 

(i) CuS0 4 + 2KCN -Cu(CN) 2 + K 2 S0 4 . 
(ii) 2Cu(CN) 2 - Cu 2 (CN) 2 + C 2 N 2 . 

Cuprous cyanide. Cyanogen. 

Further addition of potassium cyanide dissolves the cuprous 
cyanide forming a colourless solution of potassium cuprocyanide, 
K 3 Cu(CN) 4 , which ionizes into K', K', K* and Cu(CN) 4 '" and 
therefore shows none of the ionic reactions of copper. 

Cuprous sulphide, Cu 2 S, is formed when copper burns in 
the vapour of sulphur, or simply by heating sulphur and 
copper turnings together. It is a black substance, chiefly of 
historic interest, since its formation in the above way was 
observed by the Hindoos of 3,000 years ago, who therefore 
called sulphur the " enemy of copper," in Sanskrit sulvari, 
whence our sulphur. 

Cuprous hydride, said to be CuH, is precipitated as a 
yellow to brown powder on addition of a hypophosphite or 
hypophosphorous acid (p. 385) to a warm solution of copper 
sulphate. It dissolves in hydrochloric acid with evolution 
of hydrogen 

2CuH + 2HC1 = Cu 2 Cl 2 + 2H 2 . 

Cupric oxide, CuO, which is formed by strongly heating 
copper in air or oxygen, is more conveniently made by 
ignition of the nitrate 

Cu(N0 3 ) 2 = CuO + 2N0 2 + oxygen. 

It is a black hygroscopic (not deliquescent) substance, with 
the usual chemical properties of normal metallic oxides. Its 


chief use is in organic analysis, since when heated with carbon 
compounds it oxidizes their carbon to carbon dioxide and 
their hydrogen to water, both of which can easily be collected 
and weighed. 

Cupric hydroxide, Cu(OH) 2 , is precipitated in an impure 
state by adding caustic soda or potash to a solution of copper 
sulphate. It is a blue gelatinous substance, readily soluble 
in ammonia to a deep blue solution called Schweitzer's reagent. 
This liquid dissolves filter-paper, cotton-wool, and other 
forms of cellulose (C G H 10 O 5 ) n ; addition of hydrochloric acid 
reprecipitates the cellulose as a gelatinous mass which can be 
pressed through fine holes and so obtained in the delicate 
threads used in making one kind of " artificial silk." 

When a suspension of cupric hydroxide in water is boiled a 
black substance is formed, of the composition 4CuO.H 2 O 
(not CuO). This may be converted Into cupric oxide by 
strong ignition. 

Cupric chloride, CuCl 2 , may be made by dissolving the 
oxide, hydroxide or carbonate in dilute hydrochloric acid 
and evaporating the solution to crystallization. The green 
crystals of CuCl 2 .2H 2 which separate can be converted 
into the yellowish-brown anhydrous salt by gentle heat. 

When more strongly heated the anhydrous salt splits up 
into cuprous chloride and chlorine. 

The difference between the colour of the concentrated 
aqueous solution (green) and that of the dilute solution (blue) 
is elegantly explained by the ionic theory. The undissociated 
salt is yellow ; the cupric ion Cu" is blue. A concentrated 
solution, containing both undissociated molecules and ions 
will be green (yellow + blue) ; on dilution ionization rapidly 
increases and the solution turns blue since undissociated 
molecules are no longer present in appreciable numbers. If 
this is the true explanation, then addition of concentrated 
hydrochloric acid to a dilute, blue, solution should change the 
colour to green by reducing the ionization ; experiment shows 
that this happens. 

Cupric carbonate. The normal salt is unknown. The 
natural carbonates and the precipitates obtained by adding 


sodium carbonate to solutions of cupric salts are all basic 

Cupric sulphate, CuS0 4 , may be obtained as the blue 
crystalline pentahydrate, CuS0 4 .5H 2 0, by dissolving the 
oxide, hydroxide, or carbonate in dilute sulphuric acid and 
evaporating the solution until it crystallizes on cooling. It is 
manufactured by adding dilute sulphuric acid to granulated 
copper, heating to 90, and blowing air through the liquid 

Cu + H 2 S0 4 + oxygen = CuSO 4 + H 2 0. 

The process is slow, and since the copper employed generally 
contains iron ordinary commercial copper sulphate usually 
contains ferrous sulphate as an impurity. This may be 
removed by boiling a strong solution of the salt with a few 
drops of concentrated nitric acid and separating the copper 
and ferric sulphates by fractional crystallization. It is im- 
possible to separate copper and ferrous sulphates by fractional 
crystallization, since the two salts are isomorphous and form 
" mixed crystals." 

When the pentahydrated salt is heated at 100 it loses water 
and is converted into the monohydrate, CuSO 4 .H 2 0, a pale 
blue powder. At 230 the monohydrate loses practically the 
whole of its water, yielding the white anhydrous salt CuS0 4 . 
This is used as a test for water, with which it goes blue, owing 
to formation of the pentahydrate. 

Copper sulphate is a very important article of commerce. 
It is very poisonous, especially to the lower forms of plant 
life, and is therefore used as a germicide and fungicide. 
Potatoes are sprayed with " Bordeaux mixture " (copper 
sulphate and lime or limestone stirred up with water) to kill 
the fungus, Phytophthora, that causes potato-disease. Copper 
sulphate solution is also used to spray vines, wheat, etc., 
to check the development of fungi. It is used in the dye 
industry, and also in electrolytic copper-plating. 

Cupric sulphide, CuS, is obtained as a black precipitate 
when sulphuretted hydrogen is passed through a solution of 
a cupric salt acidified with hydrochloric acid. It is insoluble 
in yellow ammonium sulphide. 


Cuprammonium Compounds. The deep blue solutions 
obtained by addition of ammonia solution to copper com- 
pounds contain the copper in the form of a complex cation, 
cuprammonium, Cu(NH 3 ) 4 ". They therefore do not x give the 
reactions of the cupric ion Cu" . By adding excess of ammonia 
to copper sulphate solution and pouring a layer of alcohol 
on to the deep blue liquid, crystals of cuprammonium sulphate 
or " cuprammine sulphate" (Cu(NH 3 ) 4 S0 4 .H 2 0), may be 

Copper Alloys. 


Average Composition in parts by 


2-4 Cu. 1 Zn. 
9 Gu. 1 Sn. 
4 Cu. 1 Sn. 
3 Cu. 1 Ni. 
2 Cu. 1 Sn. 
Bronze, 99-5% ; 
55 Cu. 41 Zn. 

1 Zn. 

P, 0-5% 

Delta -metal 


Group in Periodic System : I ; Symbol : Ag ; Valency : 1 ; 
Atomic Weight: 107-88; Melting Point: 960-5; Specific 
Gravity : 10-5. 

History. Silver occurs in nature as the element and is 
easily obtained from its ores, so that it has been known for 
thousands of years. The alchemists, obsessed by astrological 
fancies, regarded it as specially connected with the moon 
and therefore called it Luna. " Luna, or Silver" says 
" GEBER," " is a Metallick Body, white with pure whitenejs, 
clean, Hard, Sounding, very durable in the Cupel, extenjible 
under the Hammer , and fujible. . . . Being put over the 
fume of acute Things, as of Vinegar, Salarmoniac, <bc., it will 
be of a wonderful Celejtine Colour [probably Geber's silver 
contained copper]. And it is a noble Body, but wants of the 
Nobility of Gold." 


Silver nitrate and chloride were both known to the chemists 
of the Middle Ages. " GEBEB " describes the preparation of 
silver nitrate very clearly 

" Dissolve Luna calcined, in Jolutive Water [nitric acid], as 
before ; which being done, coct it in a Phial with a long 
Neck, the orifice of which mujt be left unstopt, for one day 
only, until a third part of the Water be consumed. This 
being effected, Jet it with its Vefjel in a cold place, and then 
it converts to Jmall fujible Stones, like Cryftall" 

Occurrence. Native silver is found in lumps of varying 
size in South America and Mexico, etc. The chief -ores are 
horn silver (AgCl),argentite (Ag>) y pyrargyrite (3Ag 2 S.Sb a S t ), 
and strohmeyerite (Ag 2 S.Cu 2 S). 

Extraction. The methods used for the extraction of 
silver are numerous, being often conditioned by local con- 
siderations of convenience. The following are in common 

(i) Lead processes, 
(ii) Wet processes, 
(iii) Amalgamation processes. 

(i) Desilverization of Lead. If a lead ore is mixed with a 
silver ore and the mixture then smelted as for lead (p. 322), 
the lead obtained carries all the silver with it. Moreover, 
ordinary lead ores such as galena often contain silver, so that 
the desilverization of lead is of great commercial importance. 

The alloy of lead and silver can be treated in various ways 
for the extraction of the silver, according to whether it is de- 
sired to obtain both the lead and the silver or merely the silver. 

In the latter case, the alloy is cupelled, i.e. it is melted and 
strongly heated in a current of air on a cupel or vessel made 
of bone-ash [Ca 3 (P0 4 ) 2 ]. The lead oxidizes to litharge, PbO, 
part of which is carried away by the current of air while 
the rest is absorbed by the cupel. A " button " of metallic 
silver is left. This extraction of silver from argentiferous 
lead caused the alchemists to think that a partial transmu- 
tation of lead into silver had taken place. Thus ABTJ'L-QASIM 
AL- 'IRAQI (thirteenth century) says in his book Al-Muktasab, 


" Now if a part of the lead can thus be transmuted into silver, 
what is there to hinder the transmutation of the whole ? " 

If the lead-silver alloy is poor in silver, it is first concen- 
trated by removal of most of the lead, either by PATTINSON'S 
process or by PARKES' process. Pattinson's process depends 
upon the fact that when argentiferous lead is allowed to 
solidify the crystals that separate out first consist of pure 
lead ; removal of these leaves an alloy richer in silver. 
Repetition of the process results finally in the production of 
an alloy containing 2 to 3 per cent, of silver ; this is then 

Parkes' process depends upon the facts (i) that molten zinc 
will not mix with molten lead, and (ii) that silver is much more 
soluble in zinc than it is in lead. Hence, when zinc is added 
to molten argentiferous lead, most of the silver forms an alloy 
with the zinc. This alloy floats on the surface of the liquid 
and can be skimmed off ; it is then strongly heated, when the 
zinc passes off as vapour, leaving the silver behind. The 
silver still contains small quantities of lead, which are 
removed by cupellation. 

(ii) Wet processes. The chief wet process is the cyanide 
process. The silver ore is crushed and ground to a fine 
powder, which is then treated with a dilute solution of sodium 
cyanide. This dissolves the silver as sodium argentocyanide, 
NaAg(CN) 2 - 

Ag 2 S + 4NaCN = 2NaAg(CN) a + Na 2 S, 

from the solution of which metallic silver is precipitated on 
addition of zinc or in other ways. 

(iii) Amalgamation Processes. The ore is crashed, and 
mixed with about 5 per cent, of common salt. The mixture 
is placed on a stone floor and thoroughly trampled by mules, 
after which a little burnt pyrites is added, together with more 
than sufficient mercury to liberate the silver from the silver 
chloride (formed by the action of the salt upon the silver 

(i) 2NaCl + Ag 2 S = Na 2 S + 2AgCl. 
(ii) 2Hg + 2AgCl = Hg 2 Cl, + 2Ag. 


The mixture is trampled again for another six woeks, at the 
end of which time the silver is present in the mass as an 
amalgam with the excess of mercury. The amalgam is 
washed, concentrated by squeezing through canvas bags, and 
then distilled. The crude silver which is left is purified by 
cupeliation, while the mercury is condensed and used again. 

Owing to the loss of mercury (as mercurous chloride) caused 
by action (ii) above, the amalgamation process is at the present 
day obsolescent, the cyanide process gradually taking its place. 

Silver and gold are both extracted from the anode mud pro- 
duced during the electrolytic refinement of copper (p. 221). 
This is cupelled and the gold and silver may then be separated 
from one another by the action of nitric acid, which dissolves 
the silver and leaves the gold unaffected. 

Properties. Silver is a lustrous white metal which can 
take an extremely high polish. Its specific gravity is 10-5, 
its melting-point 960-5, and its boiling-point 1,955. The 
vapour of silver is blue and has a density of 54, hence its 
molecules are monatomic. Silver is unattacked by air ; the 
tarnish that appears on silver articles in towns is due to 
the presence in town air of traces of sulphuretted hydrogen, 
which acts upon silver to form black silver sulphide. 

Molten silver absorbs oxygen, which is given up when the 
metal solidifies. The bubbles of gas which come off cause a 
considerable commotion in the silver, hence the phenomenon 
is called the spitting of silver. 

Silver is practically unattacked by hydrochloric acid but 
readily dissolves in nitric acid to form silver nitrate. Cold 
dilute sulphuric acid has no action on the metal, but the 
concentrated acid dissolves it on heating, with formation of 
silver sulphate and evolution of sulphur dioxide. 

Coinage silver, formerly used for English silver coins, 
contains 92-5 per cent, silver and 7-5 per cent, copper. It is 
used for making the " silver " articles of the jeweller and 
silversmith, since the pure metal is too soft. Modern " silver " 
coins are made in England of an alloy of silver and nickel in 
approximately equal weights, with addition of a little copper. 
Recently it has become customary to add a small proportion 


of cadmium to silver used in making jewelry. The silver- 
cadmium alloy is not so difficult to manipulate as is pure silver. 
Electroplated articles are made by depositing silver electro- 
lytically upon copper ; if iron objects are to be plated they 
must first be coated with copper. The electrolytic bath used 
for the purpose of electroplating consists of a solution of 
potassium argentocyanide, KAg(CN) 2 , made by adding 
potassium cyanide to silver nitrate solution 

AgN0 3 + 2KCN = KAg(CN) a + KN0 3 . 

The object to be plated is made the cathode, while the 
anode is made of plates of pure silver. 

Sheffield Plate was made, before the invention of electro- 
plating, by welding silver on to copper and rolling the welded 
bar into sheets, from which the plated articles were made. 
This process is still employed for the reproduction of " an- 

COMPOUNDS on 1 SILVER. Silver forms one oxide, Ag 2 O, 
and one series of compounds. 1 It is uniformly univalent. 

Silver oxide, Ag 2 O, is formed as a brownish black pre- 
cipitate on adding caustic soda to a solution of silver nitrate 
2AgN0 3 + 2NaOH = 2NaN0 3 + Ag 2 + H 2 0. 

It splits up into silver and oxygen on heating, and will 
dissolve in ammonia. If the ammoniacal solution is exposed 
to air it slowly deposits a black explosive powder of fulminat- 
ing silver or silver nitride, Ag 3 N. 

Silver hydroxide, AgOH, probably does not exist in the 
free state. Moist silver oxide behaves in most reactions like 
the hydroxide might be expected to do ; thus with ethyl 
iodide, C 2 H 5 I, moist silver oxide yields ethyl alcohol and 
silver iodide 

C 2 H 5 I + " AgOH " - C 2 H 5 OH + Agl. 

Silver carbonate, Ag 2 C0 3 , is obtained as a pale yellow 
precipitate on addition of sodium carbonate to a solution of 
silver nitrate. 

Silver nitrate, AgN0 4 , or lunar caustic > can be obtained 
1 Ag a O 2 , silver peroxide, has been said to exist. 


as white tabular crystals by evaporating a solution of silver 
in nitric acid. It is used in medicine and in photography ; 
its use as an analytical reagent in the laboratory depends 
upon the fact that most silver salts are insoluble and hence 
their precipitation from a solution of silver nitrate can be 
employed as a means of recognition of the corresponding acid 
anions in the solution under examination. 

Silver nitrate is readily reduced to metallic silver in the 
presence of organic matter ; it is therefore used in solution 
as a marking-ink, since writing made with it upon linen 
rapidly turns black and the silver that is deposited cannot 
be washed out. 

On heating, silver nitrate decomposes according to the 

2AgN0 3 - 2Ag + 2NO, + O a , 

a residue of metallic silver being obtained. As the tempera- 
ture at which this decomposition occurs is much higher than 
that required to decompose copper nitrate, the latter com- 
pound may easily be separated from silver nitrate by gentle 
heat, followed by treatment with water. The water washes 
away the unchanged silver nitrate from the residual copper 

Silver chloride, AgCl, occurs naturally as born silver. It 
can be made by passing chlorine over heated silver, or by 
adding a soluble metallic chloride to a solution of silver 

e.g., NaCl + AgN0 3 = AgCl + NaNO 3 . 

Prepared in the latter way it forms a white curdy precipi- 
tate which rapidly darkens on exposure to light, finally 
becoming dark purple. 

Silver chloride is insoluble in water and in dilute acids, but 
dissolves readily in solutions of ammonia, sodium thio- 
sulphate and potassium cyanide, owing to the formation of 
soluble complex compounds 

(i) 2AgCl + 3NH a = 2AgCl . 3NH, 

A " silver ammine.*' 


(ii) Na 2 S 2 O 8 + AgCl = NaAgS 2 O 3 -f NaCL 

Sodium silver 

(iii) AgCl+ 2KCN = KAg(CN) 2 + KC1. 


Silver bromide, AgBr, and Silver iodide, Agl, are formed 
as pale yellow curdy precipitates on addition of potassium 
bromide or iodide solution to a solution of silver nitrate. 
In properties they closely resemble the chloride. Silver 
fluoride, AgF, however, is soluble in water. 

Photography. The sensitiveness of silver chloride and 
bromide to light is made use of in photography. The effect 
upon these salts of a prolonged exposure to light is obvious 
from the dark colour produced, but even an exposure of a 
fraction of a second produces a distinct effect which, while 
not apparent to the naked eye, can be rendered visible by the 
process of development. If an image is thrown, by means of 
a lens, upon a film of silver bromide, the latter is partially 
reduced to metallic silver in those parts of the film upon 
which light falls, and the extent of the reduction is propor- 
tional to the intensity of the light. If the film is now im- 
mersed in a solution of a reducing agent (" developer "), such 
as pyrogallol or hydroquinone, further reduction occurs, but 
most quickly in those parts where the reduction has been 
already started by the light. The image is thus " developed," 
but it is a negative image, since those parts which were most 
fully illuminated, on exposure, will have most silver deposited 
on them, and will therefore be blackest, on development. 
When the image has attained to a sufficient density, the 
unchanged silver bromide is dissolved out in a solution of 
sodium thiosulphate (" hypo ") and the image is thus fixed. 
Positive images can then be obtained by " printing off " the 
negative on paper covered with a film of silver chloride, and 
fixing as before. 

In making a photographic plate the silver bromide is 
precipitated in a granular form in a solution of gelatine, which 
the double purpose of holding the silver bromide and of 



rendering it more sensitive to light. The longer the gelatine 
is kept in the liquid state the larger become the granules of 
silver bromide and the greater the sensitiveness of the plate. 
This process is called ripening, and is done before the gelatine 
is spread over the plate and allowed to set. 

The plate prepared in this way is most sensitive to the 
violet and blue rays, but it may be made more sensitive to 
rays of other colours by staining the gelatine with various 
aniline dyes. 

Printing papers are covered with a film of gelatine con- 
taining a mixture of silver chloride with silver citrate. The 
latter compound increases the sensitiveness to light of the 
silver chloride. After the print has attained to a rather 
deeper shade than that required for the finished photograph, 
it is toned by means of a solution of " gold chloride," the 
action of which is to replace part of the silver in the print by 
gold, which gives a better colour. Gas-light papers are similar 
in nature to printing-out papers (except that silver bromide 
is used instead of silver chloride), but they are exposed for 
only a short time and the image is then developed as before, 
and may be toned by means of a solution of sodium sulphide 
which converts the silver into silver sulphide. 


Group in Periodic System : I ; Symbol : Au ; Valency : 
1 and 3 ; Atomic Weight : 197-2 ; Melting Point : 1,062-4 ; 
Specific Gravity : 19*43. 

History. Gold occurs native in fairly large quantities, so 
that it has been known from prehistoric times. It was sup- 
posed by the alchemists to represent perfection of the metallic 
species, and one of their aims was to convert the " baser '* 
metals into gold. Alchemical literature is full of descriptions 
of gold and of suggestions for making it. The account given 
600 years ago by the great " GEBER " is worth quoting here : 

" We say, Gold is a Metallick Body, Citrine, ponderous, 
mute, fulgid, equally digested in the Bowels of the Earth, and 
very long wajhed with Mineral Water ; under the Hammer 
extenfible, fufible, and Jujtaining the Tryal of the Cupel, and 

GOLD 235 

Cement. According to this Definition, you may conclude, 
that nothing is true Gold, unlejs it hath all the Caufes and 
differencies of the Definition of Gold. Yet, whatjoever Metal 
is radically Citrine, and brings to Equality, and cleanjeth, it 
makes Gold of every kind of Metals. Therefore, we conjider 
by the Work of Nature, and dijcern, that Copper may be 
changed into Gold by Artifice. For we Jee in Copper Mines, 
a certain Water which flows out, and carries with it thin 
Scales of Copper, which (by a continual and long continued 
Courje) it wajheth and cleanjeth. But after such Water 
ceafeth to flow, we find theje thin fcales with the dry Sand, 
in three years time to be digejted with the Heat of the Sun ; 
and among theje Scales the pure/t Gold is found. Therefore, 
We judg, thoje Scales were cleared by the benefit of the 
Water, but were equally digejted by heat of the Sun in the 
Dryness of the Sand, and Jo brought to Equality. Wherefore, 
imitating Nature, as far as we can, we likewije alter ; yet in 
this we cannot follow Nature. 

" Aljo Gold is of Metals the mojt precious and it is the 
Tincture of Eednejs ; becauje it tingeth and transforms every 
Body. It is calcined and dijjolved without profit, and is a 
Medicine rejoycing, and conjerving the Body in Youth. It 
is most eajily broken with Mercury, and by the Odour [i.e. 
vapour] of Lead. There is not any Body that in act more 
agrees with it in Subjtance than Jupiter [tin] and Luna 
[silver] : but in Weight, Denfeness, and Putrefcibility, Saturn 
[lead], in Colour Venus [copper] ; in Potency indeed Venus 
is more, next Luna, then Jupiter, and then Saturn : but la/tly 
Mars [iron]. And this is one of the Secrets of Nature. Like- 
wije Spirits are commixed with it, and by it fixed, but not 
without very great Ingenuity, which comes not to an Artificer 
of a Jtiff neck." 

Occurrence. Native gold is found in quartz veins and, 
in smaller quantities, in alluvial deposits. The lumps vary 
in size from huge nuggets (the largest yet recorded weighed 
nearly two hundredweight !) to minute grains. Gold is widely 
distributed in the earth, but in only a few districts is enough 
found to make the extraction worth while. Sea-water con- 


tains minute traces, but, in spite of many Attempts, no 
commercially successful method of extracting gold from sea- 
water has been invented. At the present price of gold, a ton 
of sea-water contains about sevenpenny worth. The chief 
gold-producing regions are the Transvaal, Australia, North 
America (Klondike, etc.), and Russia. 

Extraction. Nuggets need no extraction. Gold-bearing 
alluvial sand is treated by levigation, that is, it is stirred in a 
current of water which washes away the lighter particles 
while the heavy grains of gold fall to the bottom. A better 
process is to treat the auriferous sand or crushed ore with a 
dilute solution of sodium cyanide (this is where most of the 
sodium manufactured is used), which dissolves the gold in the 
form of sodium aurocyanrh, NaAu(CN) 2 . The gold may be 
recovered from this solution by electrolysis, or by precipita- 
tion with metallic zinc, which replaces the gold in the complex 
cyanide. [N.B. The cyanide employed is called " potas- 
sium " cyanide, but is really sodium cyanide, as stated.] 

Gold is also extracted by a process of amalgamation. The 
alluvial sand or finely stamped ore is mixed with mercury, 
which forms an amalgam with the gold. The pulp or slime 
is then run down a table covered with amalgamated copper 
plates, the length of the table being some 20 feet and its 
gradient 1 or 2 inches per foot. The greater portion of the 
gold is retained on the plates, whence the amalgam is scraped 
off. The latter is squeezed in chamois-leather bags, to remove 
excess of mercury, and the residue is then distilled in iron 
retorts. The gold left in the retorts is purified by cupellation. 
Any gold not removed from the slime by the amalgamation 
process is extracted by the cyanide method. 

Gold is refined at the present day chiefly by electrolysis. 
The crude gold (generally containing silver) is made the 
anode in an electrolytic cell containing a solution of gold 
chloride acidified with hydrochloric acid. A thin sheet of 
pure gold is used as cathode, and when the current is passed 
the anode decreases in bulk while the cathode increases. As 
a matter of practical importance, it is found advisable to 
superimpose an alternating current upon the direct current, 

GOLD 237 

since this prevents accumulation of solid silver chloride upon 
the anode. 

Properties. Gold is a yellow metal with a red reflex. 
It has a specific gravity of 1943, and it melts at 1,062-4 to 
a green liquid. It is more malleable and ductile than any 
other metal, and is unaffected by air, oxygen, and acids 
(except aqua regia). It is attacked by chlorine, forming gold 
chloride, and by potassium cyanide in the presence of oxygen, 
forming the aurocyanide 

4Au + 8KCN + 2H 2 O + 2 = 4KAu(CN) f + 4KOH. 

It dissolves in mercury to form an amalgam, and in aqua 
regia to form " gold chloride." (See below.) 

The preparation of colloidal gold, which is carried out by 
reducing a solution of gold chloride with formaldehyde (or in 
other ways), is of interest as representing the realization of 
a dream of alchemy the making of " potable gold," or gold 
in a drinkable form. The " solutions of gold " prepared by 
the alchemists were solutions of soluble gold compounds. 
MOSES' chemical experiments on the golden calf possibly 
consisted in heating the calf with liver of sulphur and thus 
converting it into the soluble salt NaAuS, which could then 
have been dissolved in water for the Israelites to drink. 

CoUoidal gold made in the above way is red, but a blue 
form is also known ; this is made by using hydrazine hydrate 
instead of formaldehyde as the reducing agent. 

Purple of Cassius (colloidal stannic acid containing colloidal 
gold) is obtained as a purple-violet precipitate by adding a 
mixture of stannous and stannic chlorides to a dilute solution 
of gold chloride. It is used for giving a red colour to glass 
and porcelain and is caUed after its discoverer, CASSIUS (1685). 

Gold is chiefly used for coinage and for jewelry. It is too 
soft to be used in the pure state and therefore is alloyed with 
copper or silver. The proportion of gold in the alloy is 
expresssed in carats, pure gold being taken as of 24-carat 
purity. Thus 18-carat gold is a copper-gold alloy containing 
J-f of its weight of gold. The English sovereign used to be 
22-carat. The fineness of a gold alloy is roughly ascertained 


by the touchstone, a siliceous mineral. The alloy is rubbed 
on the stone and leaves a streak of metal behind ; this is 
treated with nitric acid and its behaviour compared with that 
of another streak made with an alloy of known fineness. 

COMPOUNDS OF GOLD. The compound formed when gold 
is dissolved in aqua regia is not really gold chloride but 
hydrochloro-auric acid, HAuCl 4 . This may be obtained in the 
form of trihydrated yellow crystals, HAuCl 4 .3H 2 0, on 
evaporation. The acid itself and its sodium salt are used in 
photography, under the general name of " chloride of gold," 
for toning silver chloride prints. The action is due to the 
reduction of the " gold chloride " and deposition of metallic 

Gold forms two series of compounds, the aurous and auric. 
In the first it is univalent and in the second tervalent. Gold 
compounds are as a rule unstable and yield metallic gold on 
heating. Auric oxide and hydrochloro-auric acid appear to 
have been known to the alchemists. Fulminating gold ia 
made by adding ammonia to gold chloride solution ; it is a 
greenish powder that explodes violently when struck. 


1. Give an account of the chemical history of copper, silver and 
gold. To what chemical reaction do you suppose the alchemists 
referred in the sentence " Sol is devoured by the Red Dragon " ? 

2. Compare the properties of copper with those of (a) silver, (6) 
nickel (p. 560). 

3. Estimate the success or failure of the periodic system in grouping 
the alkali metals with copper, silver and gold. 

4. Give an account of the metallurgy of copper. 

5. Write a short account of the processes employed in the extraction 
of silver from its ores. 

4. Explain the chemistry of photography as far as you can. 


TYPICAL ELEMENTS : Beryllium, Magnesium. 
Sub-group A (similar to the typical elements) : Calcium, 

Strontium, Barium, Radium. 
Sub-Group B : Zinc, Cadmium, Mercury. 


The elements magnesium, calcium, strontium and barium 
are the alkaline earth metals. They are all comparatively 
" light " metals, since the specific gravity of the heaviest of 
them, barium, is only 3*6. Their valency is uniformly 2. 

All of them show considerable chemical activity, and 
exhibit a close resemblance to one another in their compounds. 

Magnesium has many similarities to zinc, as well as to 
calcium, strontium and barium. 


Group in Periodic System : II ; Symbol : Mg ; Valency : 
2 ; Atomic Weight : 24-32 ; Melting Point : 650 ; Specific 
Gravity : 1-74. 

History. Magnesium carbonate was known to the ancient 
Assyrians, since a votive tablet from the palace of King Sargon 
was found by BERTHELOT to consist of this substance. The 
alchemists gave the name magnesia to very many different 
alloys and oxides, and became hopelessly confused when 
trying to describe its properties, as might have been expected. 
They seem, however, to have distinguished between a black 



magnesia, which was possibly pyrolusite (manganese dioxide) 
and a white magnesia, which may have been magnesium 
carbonate, although this substance is comparatively rare in 
nature in the pure state. 

Magnesium sulphate was obtained in 1695 by GREW, from 
the water of a spring at Epsom, and by addition of sodium 
carbonate to a solution of this " Epsom salt " a white pre- 
cipitate of magnesium carbonate was formed. This was 
called magnesia alba and was shown by BLACK (1755) to be 
a compound of fixed air (carbon dioxide) with a peculiar 
" earth " magnesia (MgO). In 1800 DAVY succeeded in 
isolating a metal, magnesium, from magnesia ; his product 
was, however, impure and magnesium was prepared fairly 
pure only some thirty years later by BUSSY. 

Occurrence. Magnesium occurs chiefly as its carbonate, 
MgC0 3 . When it occurs singly this compound is caUed 
magnesite, but it is found more often combined with 
calcium carbonate in dolomite, MgCO 3 .CaC0 3 , of which the 
Dolomitic Alps are largely composed. Magnesium occurs 
also as chloride in the mineral carnallite, KCl.MgCl 2 .6H 2 
(Stassfurt) ; as silicate in talc, Mg 3 H 2 (SiO 3 ) 4 , and as- 
bestos, CaSi0 3 .3MgSi0 3 ; and as sulphate in kainite, 
K 2 S0 4 .MgSO 4 .MgCl 2 .6H 2 and kieserite, MgS0 4 .H 2 0. It 
has been shown by WILLSTATTER that the molecule of chloro- 
phyll, the green colouring-matter of plants, contains an atom 
of magnesium. 

Extraction. Magnesium is manufactured from the car- 
nallite mined at Stassfurt. When this substance is heated it 
loses its water of crystallization and then melts to a colourless 
liquid consisting of fused potassium and magnesium chlorides. 
This is electrolysed in an iron pot, which is made the cathode 
the anode being a carbon rod which dips into the molten mass. 
In practice, a little calcium fluoride is added to the fused 
electrolyte. Chlorine is evolved from the anode and mag- 
nesium is liberated at the cathode. The temperature of the 
fused chloride is above the melting-point of the magnesium, 
and the latter therefore melts and rises to the surface, whence 
it is run off. Owing to the easy oxidation of magnesium, the 


air in the apparatus is replaced by an inert gas such as coal-gas 
(not nitrogen, since magnesium combines with nitrogen to form 
magnesium nitride). 

Properties. Magnesium is a white silvery metal which is 
stable in dry air, but slowly oxidizes in moist air, forming mag- 
nesium oxide, MgO. It has a specific gravity of 1-74, melts 
at 650, and boils at 1,100. It is ductile and malleable and 
is usually met with as magnesium ribbon or wire. In making 
magnesium wire the metal has to be heated ; ribbon is made 
from the wire by means of heavy rollers. 

When heated in the air magnesium takes fire and burns 
with a brilliant white light rich in the chemically active rays. 
It is therefore used in photography, star-shells, etc. The 
product of the combustion of magnesium in air is a mixture of 
the oxide MgO and nitride Mg 3 N 2 . Magnesium will burn in 
steam (p. 181) and also in carbon dioxide ; in each case the 
oxide of the metal is formed and the other element (hydrogen 
or carbon) liberated. 

Magnesium is employed commercially in the preparation 
of " flashlight powders " (KC10 8 + Mg powder, etc.), fire- 
works, and the alloy (with aluminium) called magnalium, 
which is very light but surprisingly strong and has therefore 
been used in air ship- construction, balance-making, and so on. 

Magnesium is readily soluble in dilute acids, but caustic 
alkalis have no action on it. It dissolves in dry ethereal 
solutions of alkyl halides, etc., forming " GRIGNARD com- 
pounds," which are of great importance in organic chemistry 

CH 3 I + Mg - CH 3 - Mg - I. 

Methyl iodide. Methyl magnesium iodide, 

a " Grignard compound." 


Magnesium oxide or magnesia, MgO, is obtained 
(i) when magnesium burns in air or oxygen ; 
(ii) by heating the carbonate 

MgCO.^MgO + CO,; 
(iii) by heating the nitrate 

2Mg(N0 3 ) a = 2MgO + 4NO a + 0, ; 


(iv) by heating the hydroxide 

Mg(OH) 2 = MgO + H 2 0. 

It is a white powder, very slightly soluble in water, with 
which it slowly combines, forming the hydroxide 

MgO + H 2 ^= Mg(OH) 2 ;= Mg" + 20H'. 

The solution has an alkaline reaction. It is readily soluble 
in acids, yielding magnesium salts. It is employed in 
medicine and also as a fire -resisting or refractory material for 
electric furnaces. NEKNST electrical lamps, formerly much 
employed but now obsolescent, consist of a filament of mag- 
nesium oxide (mixed with oxides of other metals) ; on heating 
this filament it becomes a conductor and is raised to a high 
temperature on passage of the electric current, giving out a 
very bright white light. 

Magnesium hydroxide, Mg(OH) 2 , is obtained as a white 
precipitate by addition of caustic soda, caustic potash or 
ammonium hydroxide solution to a solution of a magnesium 

2NaOH + MgS0 4 = Na 2 SO 4 + Mg(OH) 2 . 

If ammonium chloride is added to a solution of a mag- 
nesium salt, and ammonium hydroxide added afterwards, no 
precipitate is obtained. This is easily explained by the ionic 
theory (see p. 135). The ammonium chloride suppresses the 
ionization of the ammonium hydroxide so much that the 
solubility-product of magnesium hydroxide is not reached, 
and therefore no precipitation occurs. For the same reason, 
magnesium hydroxide is dissolved by a solution of ammonium 

Magnesium peroxide, Mg0 2 , is obtained in an impure 
state by adding sodium peroxide to magnesium sulphate 
solution. Ifc is a white powder with oxidizing powers, and 
is used as an antiseptic. 

Magnesium sulphate, MgS0 4 , may be made by dissolving 
the metal, oxide, or carbonate in dilute sulphuric acid and 
evaporating the solution to crystallization, when colourless 
crystals of magnesium sulphate heptahydrate, MgS0 4 .7H a O t 


or Epsom salt, separate. On heating, these crystals lose 
water, forming lower hydrates and finally the anhydrous salt. 

It readily forms double salts with the sulphates of sodium 
and potassium ; these double salts have the general formula 
M 2 S0 4 .MgS0 4 .6H 2 0, and in solution give all the ionic 
reactions that would be shown by the two sulphates of which 
they are composed. 

" Epsom salt " is used as a purgative and in the dye 
industry. It is sometimes used as an artificial manure for 
certain soils. 

Magnesium chloride, MgCl 2 . The deliquescent crystals 
obtained from the solution made by dissolving magnesium or 
its oxide or carbonate in dilute hydrochloric acid have the 
composition MgCl 2 .6H 2 0. On heating, this hexahydrated 
salt loses water and hydrochloric acid, forming an oxychloride, 

Mg<^ which itself splits up at a higher temperature into 

magnesium oxide and hydrochloric acid. 

Anhydrous magnesium chloride is therefore best prepared 
by heating magnesium in a stream of hydrochloric acid gas, 
or by driving off the water and ammonium chloride from the 
double salt NH 4 Cl.MgCl 2 .6H 2 by heat 

NH 4 Cl.MgCl 2 .6H 2 = NH 4 C1 + 6H 2 + MgCl 2 . 

Magnesium carbonate. When sodium carbonate is 
added to a solution of a magnesium salt the precipitate 
consists not of normal magnesium carbonate but of a basic 
carbonate. This may be converted into the normal carbonate 
by suspending it in water and blowing a current of carbon 
dioxide through. Excess of carbon dioxide causes the 
carbonate to dissolve as bicarbonate 

MgC0 8 (insoluble) + H 2 O + C0 2 ^= Mg(HC0 3 ) 2 (soluble). 


Group in Periodic System : II ; Symbol : Ca ; Vakncy : 2 5 
Atomic Weight : 40-07 ; Melting Point 810 ; Specific 
Gravity : 1-55. 


History. LAVOISIER suggested that lime was the oxide 
of a metallic element, and in 1808 DAVY prepared an impure 
specimen of calcium by electrolysis. A purer specimen was 
made by BUNSEN in 1855, and in 1898 MOISSAN obtained the 
metal by strongly heating calcium iodide with excess of 
sodium. Compounds of calcium, such as quicklime, slaked 
lime, and limestone have of course been known and used for 
thousands of years, since many of them occur naturally in 
large quantities. 

Occurrence. Calcium carbonate, CaC0 3 , is the chief 
calcium mineral. This is found in various crystalline forms, 
calcite, aragonite, marble, Iceland spar, etc., and in the amor- 
phous state as limestone and chalk. Dolomite is a double 
carbonate of calcium and magnesium, CaCO 3 .MgC0 3 . 

Calcium is also found in the form of its sulphate ; gypsum, 
selenite, and alabaster are CaSO 4 .2H 2 O and anhydrite is 
CaS0 4 . Phosphates of calcium occur n,aturally (see phos- 
phoruSj p. 373). 

Calcium is widely distributed in the soil and is found in all 
plants and animals. Eggshells consist largely of calcium 

Preparation. Calcium is becoming of importance in 
technical chemistry and is therefore prepared on the large 
scale. In RUFB and PLATO'S method a mixture of 100 parts 
of anhydrous calcium chloride and 16 parts of calcium fluoride 
(fluorspar) is fused at a temperature of 660 and electrolysed 
in a pot lined with graphite. The graphite is made the anode 
and the cathode is a vertical iron wire or rod upon which 
the calcium is deposited during electrolysis. The calcium 
obtained in this way is very pure. 

Properties. Calcium is a hard silver- white metal with 
a slight tinge of yellow. It oxidizes slowly in moist air and 
attacks water in the cold, with evolution of hydrogen 
Ca + 2H*0 = Ca(OH) 2 + H 2 . 

It has a specific gravity of 155 and melts at 810 C. If 
heated in air or oxygen it burns with a reddish flame, forming 
calcium oxide, CaO. It will also combine with nitrogen if 


heated in this gas, being converted into calcium nitride, 
Ca 3 N 2 . It is a good reducing agent. Metallic calcium is 
used in the laboratory to free alcohol from traces of water. 

COMPOUNDS OF CALCIUM. Calcium hydride, CaH 2 , is 
obtained when calcium is heated in hydrogen. 

Ca + H 2 = CaH 2 . 

It is a colourless substance, used commercially, under the 
name of hydrolith, for the preparation of hydrogen for air- 
ships, etc. 

CaH 2 + 2H 2 = Ca(OH) 2 + 2H 2 . 

Calcium fluoride, CaF 2 , occurs naturally as fluorspar. 

Calcium chloride, CaCl 2 , is the chief waste-product in the 
SOLVAY process for the manufacture of sodium carbonate. In 
the laboratory it may be made by dissolving the oxide or 
carbonate of the metal in dilute hydrochloric acid. On 
evaporation of the solution calcium chloride hexahydrate, 
CaCl 2 .6H 2 0, separates in the form of colourless and extremely 
deliquescent crystals. On heating these crystals they are 
converted into the dihydrate, CaCl 2 .2H 2 0; further heating 
results in formation of the anhydrous salt, CaCl 2 , which is 
largely used in the laboratory as a drying agent, since it has 
a great affinity for water and is not very reactive. It cannot, 
however, be used for drying ammonia (p. 338) as it absorbs 
this gas, forming compounds which have the general formula 
a;CaCl 2 .2/NH 3 ; neither must it be used for drying alcoholic 
solutions, since it dissolves in alcohol. It is also unsuitable 
as a dehydrating agent for aniline and amines, in general. 

Calcium oxide or quicklime, CaO, is formed when the 
metal burns in air and when the nitrate or carbonate is strongly 
heated. Commercially it is made by heating limestone in 
a lime-kiln. The crushed limestone is mixed with coal or 
coke and the kiln filled with the mixture which is then lit 
at the bottom. The lime, in the form of a powder or small 
lumps, is removed from the bottom of the kiln and fresh 
mixture added at the top. Lime prepared in this way con- 
tains the ashes of the coal or coke used, but these are rarely 


harmful, and are often beneficial in lime which is used tot 
commercial purposes. 

Calcium oxide is a white amorphous substance very difficult 
to fuse. It melts in the electric furnace and can even be 
boiled ; it is not, however, melted by the oxyhydrogen flame 
but merely made white hot. White-hot quicklime gives a 
brilliant light, called the limelight. Owing to its refractory 
nature lime is often used in the construction of electric 

When water is added to quicklime, evolution of heat occurs 
and calcium hydroxide or slaked lime, Ca(OH) 2 , is 
formed. If the quicklime is in lumps, these will fall to pieces 
and the slaked lime will be obtained as a powder. Calcium 
hydroxide is slightly soluble in water ; the solution is called 
lime-water, and is used as a test for carbon dioxide and as a 
mouth- wash for infants. If excess of lime is stirred up with 
water the mixture is called milk of lime. Soda-lime is made 
by slaking quicklime with caustic soda solution ; it behaves 
like caustic soda in reactions, but has the advantage of being 
non-deliquescent and not easily fusible, and may therefore 
be often conveniently employed instead of solid caustic soda 
(e.g., in the preparation of methane from sodium acetate, 
p. 294). 

Lime finds many applications both in the laboratory and 
elsewhere. Technically it is employed in the purification of 
coal-gas, in the manufacture of caustic soda and sodium 
carbonate, in tanning, in glass-making, in the manufacture 
of bleaching-powder, in the refinement of sugar, and for many 
other purposes. It is also the main constituent of mortar 
and cement. 

Mortar consists of a pasty mixture of slaked lime, sand, 
and water. The " setting " of mortar is due chiefly to loss 
of moisture by evaporation, but the lime slowly changes into 
caicium carbonate by the action of atmospheric carbon 
dioxide, and part of it combines with the sand to form cal- 
cium silicate. 

Hydraulic cement is so called because it will set even under 
water. To make it, limestone containing a little clay IB 


heated in the lime-kiln, and the quicklime BO obtained slaked 
and mixed with sand and water. It is very useful for build- 
ing-work which has to stand in water. 

Portland cement is made by strongly heating a mixture of 
limestone, or chalk, and clay. The product is then finely 
powdered and is ready for use. When mixed thoroughly 
with water it quickly sets to a hard mass. 

Concrete is a mixture of Portland cement and small gravel 
or finely broken bricks. In reinforced concrete the concrete 
is allowed to set over a skeleton of steel girders or rods. 

Calcium carbonate, CaC0 3 , occurs naturally in many 
forms, for which see p. 244. It can be obtained as a white 
precipitate by passing carbon dioxide through lime-water 

Ca(OH) 2 + C0 a = CaC0 3 + H 2 0. 

It is insoluble in water, but dissolves in an aqueous solution 
of carbon dioxide, owing to the formation of the soluble 
bicarbonate, Ca(HCO 3 ) 2 

CaC0 3 + H 2 O + CO 2 = Ca(HC0 3 ) a . 

On boiling a solution of the bicarbonate the normal carbonate 
is re-formed and is precipitated. For the effect of the pre- 
sence of calcium bicarbonate in water see p. 445, Hardness 
of Water. 

Calcium sulphate, CaS0 4 . The naturally occurring forms 
of this substance have already been mentioned, p. 244. When 
gypsum, CaSO 4 .2H 2 0, is carefully heated to 120-130, it 
is converted into the hemihydrate, CaS0 4 .iH 2 0. This is 
known as Plaster of Paris ; when it is made into a thick paste 
with water it rapidly sets to a mass of gypsum, and as expan- 
sion takes place in the process, sharp casts may be obtained 
by allowing the mixture to set in moulds. 

If the gypsum is heated to a high temperature it loses all 
its water of crystallization and is converted into the anhy- 
drous salt, CaS0 4 . This will not take up water, and is said 
to be " dead-burnt." 

Calcium bisulphite, Ca(HS0 8 ) 2 . See p. 466. 


Calcium nitrate, Ca(N0 3 ) 2 , and basic calcium nitrate. 
See p. 355. Calcium nitrate forms colourless deliquescent 
crystals with four molecules of water of crystallization, 
Ca(N0 3 ) 2 .4H 2 0. When strongly heated it splits up into 
quicklime, oxygen, and nitrogen peroxide 

2Ca(N0 3 ) 2 = 2CaO + 4N0 2 + O,. 

Calcium carbide, CaC 2 , is made by heating a mixture of 
quicklime and powdered coke in the electric furnace. It is 
decomposed by water, with evolution of acetylene 

CaC 2 + 2H 2 = Ca(OH) 2 + C 2 H 2 . 

When calcium carbide is strongly heated in a current of 
nitrogen calcium cyanamide, CaCN 2 , is formed 

CaC 2 + N 2 = CaCN 2 + C. 

Calcium cyanamide is used as an artificial manure 
(pp. 342, 371). 

Calcium phosphate, Ca 3 (P0 4 ) 2 . See p. 373. 

Calcium "superphosphate," CaH 4 (P0 4 ) 2 , is manufactured 
in enormous quantities for use as a fertilizer in agriculture. 
It is made by adding to crushed calcium phosphate the 
exact weight of sulphuric acid required according to the 

Ca 8 (P0 4 ) 2 +2H 2 S0 4 +4H 2 0-2CaS0 4 .2H 2 0+CaH 4 (P0 4 ) 2 . 

Commercial " superphosphate " is therefore a mixture of 
calcium tetrahydrogen phosphate, CaH 4 (P0 4 ) 2 , with gypsum. 
Bleaching -powder, Ca(OCl)Cl. Although absolutely dry 
chlorine will not act upon specially dried calcium hydroxide, 
ordinary " dry " slaked lime will readily absorb chlorine, 
forming a white powder called bleaching-powder. On the 
commercial scale bleaching-powder is made by passing a 
stream of chlorine slowly over a layer of slaked lime on the 
floor of a series of chambers made of stone or lined with 
lead. The reaction that occurs may be represented by the 

(CaOH) a + C1 2 = CaOCl 2 + H 2 0. 


Bleaching- powder is, as its name implies, used for bleaching, 
since with a dilute acid it reacts to yield free chlorine 
CaOCl 2 + H 2 S0 4 = CaS0 4 + H 2 + C1 2 . 

Constitution of Bleaching -powder. The constitution of 
bleaching-powder proved a very difficult matter to settle. 
GAY-LUSSAO considered it to be a loose compound of lime 
and chlorine, "chloride of lime," CaO.Cl 2 , whereas BALARD 
regarded it as a mixture of calcium chloride and calcium 
hypochlorite, CaCl 2 .Ca(OCl) a . Against Balard's theory there 
are the facts that 

(i) Bleaching-powder is not deliquescent, whereas calcium 
chloride is. 

(ii) Calcium chloride is soluble in alcohol, but cannot be 
dissolved out of bleaching-powder by this solvent. 

(iii) Carbon dioxide will act upon bleaching-powder with 
liberation of practically all the chlorine from the latter 
compound ; carbon dioxide, however, has no action on 
calcium chloride. 

(iv) Bleaching-powder cannot be made by mixing equi- 
molecular proportions of calcium hypochlorite and calcium 

ODLING therefore in 1861 suggested that bleaching-powder 

was chiefly calcium chlorohypochlorite, Ca<^ , that is, a 


" mixed salt," half way between calcium chloride, Ca<^ , and 

X C1 
calcium hypochlorite, Ca/ . This formula agrees with all 

X OC1 

the reactions of the substance, the formula for which, allowing 
for the water always present, may be written Ca(OCl)Cl.H 2 O. 
According to this formula 1 the " available " chlorine (i.e. the 


chlorine liberated from it by a dilute acid) should be - , or 


1 Olding's formula does not account for the existence of a lithium 
bleaching-powder, but i* the best formula yet suggested. 


about 49 per cent., of the weight of bleaching-powder taken. 
The average percentage of available chlorine is * consider ably 
less, generally 35-38 per cent., and diminishes on keeping the 
powder, possibly owing to the change 

6CaOCl a = 5CaCl a + Ca(C10 3 ) 2 . 

Calcium chlorate. 

Commercial bleaching-powder is therefore bought and sold 
on the basis of its available chlorine content. This may be 
estimated by grinding up a known weight of the powder with 
water in a mortar and making up the fine paste to a known 
volume with distilled water in a graduated flask. An aliquot 
portion of the well-shaken suspension is then mixed with 
excess of potassium iodide solution, acidified with dilute 
hydrochloric acid, and titrated with standard sodium 

(i) CaOCl 2 + 2HC1 = CaCl 2 + H 2 O + Cl* 
(ii) C1 2 + 2KI = 2KC1 + I 2 . 
(iii) I t -f 2Na 2 S 2 3 = 2NaI + Na 2 S 4 6 . 

Sodium tetrathionate. 

1 litre N/10 sodium thiosulphate =s 12-7 gms. iodine == 3-55 
gms. chlorine. For other methods, see BERRY'S Volumetric 

WILKS and others have shown that bromine and iodine 
probably form similar compounds with slaked lime, brom- 
bleaching pounder and iodine bleaching-powder, CaOBr 2 and 
CaOI 2 , thus confirming previous observations by BERZELIUS. 


Group in Periodic System : II ; Symbol : Sr ; Valency : 2 ; 
Atomic Weight: 87-63; Melting Point: 800; Specific 
Gravity: 2-55. 

History. In 1790 CRAWFORD concluded that the mineral 
strontianite (SrC0 8 ), so called because it is found near 
Strontian in Argyllshire, contained a new earth ; in 1791 
HOPE and in 1792 KLAPROTH confirmed Crawford's con- 
clusion, and the metal strontium was isolated sixteen veara 
later by DAVY, who electrolysed the fused chloride. 


Occurrence. The chief ores of strontium are strontianite, 
SrC0 3 , and celestine, SrS0 4 . These are fairly widely distri- 
buted, but usually in small quantities. 

Preparation. The metal may be made by Davy's method 
or by the thermite process (see p. 277). 

Properties. Strontium is very similar to calcium in 
physical and chemical properties. It melts at 800 and has 
a specific gravity of 2-55. The oxide, SrO, and hydroxide, 
Sr(OH) 2 , are used in sugar-refining, since they combine with 
sugar to form a white insoluble substance, C 12 H 22 O n .2SrO, 
called strontium saccharate. This is washed free from im- 
purities, stirred up in water, and decomposed by a stream of 
carbon dioxide, which liberates the sugar and precipitates 
the strontium as carbonate 

C 12 H 22 O n .2SrO + 2CO, = C 12 H 22 O n + 2SrCO,. 


The insoluble carbonate is allowed to settle and the clear 
solution of pure sugar evaporated to crystallization. By 
strongly heating the carbonate it is reconverted into the 
oxide, available for further use. 

Strontium nitrate, Sr(N0 3 ) 2 , is used in pyrotechny for 
making red flares. 


Group in Periodic System : II ; Symbol : Ba ; Valency : 2 j 
Atomic Weight : 137-37 ; Melting Point : 850 ; Specific 
Gravity: 3-6. 

History. The first mention of a barium compound seems 
to be in 1602, when a shoemaker of Bologna, CASCIOBOLTJS, 
drew attention to the fact that if a mineral called heavy-spar 
or barytes (BaS0 4 ) was heated strongly with a combustible 
substance, such as charcoal, the product possessed the power 
of shining in the dark, after exposure to sunlight. This 
phosphorescent substance (BaS) was called the Bolognan Stone, 
and when phosphorus was discovered nearly a century later 
the two were at first considered to be in some way related. 

SCHEELE, however, in 1774, and GAHN, showed that heavy. 


spar was the sulphate of an " earth " or metallic oxide, and 
DAVY succeeded in isolating the metal, in an impure state, in 
1808, by electrolysis of the fused chloride. 

Occurrence. Barium occurs chiefly as the sulphate, 
heavy-spar or barytes, BaS0 4 , and also as the carbonate, 
witherite, BaC0 8 . 

Preparation. Barium can be made by Davy's method 
(above), or by the thermite process (see p. 277). 

Properties. In appearance and general chemical pro- 
perties, barium closely resembles calcium and strontium. Its 
specific gravity is 3-6 and its melting-point 850. Barium 
oxide, BaO, was formerly used in the manufacture of oxygen 
by Erin's process : on heating in air it is converted into the 
peroxide, Ba0 2 , which is reconverted into the normal oxide, 
with loss of oxygen, on further heating or on reducing the 

2BaO + 2 ;=2Ba0 2 . 

Barium chloride, BaCl 2 .2H 2 0, is used in solution as a 
reagent in analysis, especially in testing for sulphates, with 
which it gives a white insoluble precipitate of the sulphate, 
BaS0 4 . 

Barium nitrate, Ba(N0 8 ) 2 , is used in the production of green 
flares, in the firework industry. 

Barium sulphate, BaS0 4 , is one of the most insoluble salts 
known. It is used as a white pigment, under the name of 
" permanent white." 


1. Give an account of the occurrence, extraction and uses of mag- 

2. Compare the chemical properties of magnesium with those of (a) 
calcium, and (6) zinc. 

3. State the chemical nature of the following : alabaster, plaster of 
Paris, mortar, gypsum, concrete, fluorspar, Iceland spar. 

4. Discuss the constitution of bleaching-powder. 

& What do you know of the history of the alkaline-earth metal* f 


GROUP II, Sub-group B 



Oroup in Periodic System : II ; Symbol : Zn ; Valency : 
2 ; Atomic Weight : 65-38 ; Melting Point : 4194 ; Specific 
Gravity: 6-9. 

History. In the form of its alloy with copper, brass, zinc 
was known to the chemists of antiquity, who were also 
acquainted with its oxide and carbonate, which they 
grouped together under the name of cadmia, afterwards 
changed to tutia. Brass was made by heating copper with 
tutia, as described by many alchemists. The name zinc seems 
to have been given to the metal by PARACELSUS (1493-1541), 
but it was not till 1720 that zinc was obtained in a state of 
approximate purity by HENCKEL. HOMBERG, however, had 
shown in 1695 that zinc could be extracted from zinc blende, 

Occurrence. Zinc does not occur native. Its chief ores 
are zinc blende (ZnS), and calamine, ZnC0 8 . It also 
occurs as zinc ferrite (Franklinite), Zn(Fe0 2 ) 2 , in the United 
States. The principal deposits of zinc ores are found 
in Great Britain, Belgium, Germany, Australia and North 

Extraction. The manufacture of zinc was begun in 
Bristol in 1743, and was practically a Bristol monopoly until 
1807, when a zinc works was erected at Liege, followed 




by many more in Belgium and in Silesia. The Belgian and 
Silesian processes differ only in detail. 

The ore is strongly heated in air, when (i) zinc blende is 
oxidized and (ii) calamine loses carbon dioxide, the product 
in each case being zinc oxide. 

(i) 2ZnS + 30 2 = 2ZnO + 2S0 2 . 
(ii) ZnCO, = ZnO 4- CO.. 

The sulphur dioxide given oif when blende is used is con- 


Fia. 55. Belgian Zinc Plant. 

verted into sulphuric acid and is an important factor in the 
commercial success of the operation. 

The zinc oxide obtained in the first operation is then mixed 
with powdered coal and strongly heated in fireclay retorts by 
means of producer-gas (p. 301). The Belgian and Silesian 
processes differ in the shape of the retorts and receivers and 
in the arrangement of the furnaces ; the Silesian method 
is the better of the two for ores that are not rich in 



The zinc oxide is reduced in the retorts, and zinc vapour 
comes over 

ZnO + C = Zn + CO. 

Part of the zinc collects as a fine powder (" zinc-dust ") and 
part condenses to a liquid which is allowed to solidify in 
moulds. The crude zinc contains many other metals as 
impurities, chiefly lead, antimony and arsenic ; it is purified 
by heating the molten metal in a reverberatory furnace for a 
considerable time, when certain of the impurities rise to the 
surface and can be removed. Electrolytic refinement has 
recently been introduced, with much success, the metal pre- 




Fia. 56. Silesian Zinc Plant. 

pared in this way having a purity of approximately 99-95 per 
cent. By the electrolysis of specially purified zinc chloride, 
zinc of 99-998 per cent, purity may be obtained without 
difficulty. Commercial zinc (known as spelter) always con- 
tains lead (0-02-1-5 per cent.) and iron (0-01-0*05 per cent.). 
A small percentage of lead is advantageous if the metal has 
to be rolled. 

Properties. Zinc is a hard and rather brittle bluish- white 
metal ; it becomes ductile and malleable when heated from 
100 to 150, but becomes brittle again at 200, at which tern- 
perature it can be powdered. Its specific gravity is 6'9, 


melting-point 419-4, and boiling-point 920. Zinc is stable 
in dry air, but in moist air is slowly converted into a white 
" rust," basic zinc carbonate. If heated strongly in the air 
zinc burns with a greenish blue flame, forming clouds of the 
very light zinc oxide (philosopher's wool). 

The general chemical properties of zinc are well known ; 
the metal readily dissolves in acids, and is soluble also in hot 
aqueous caustic alkalis 

Zn + 2NaOH = Na 2 ZnO 2 + H 2 . 

Sodium zincate. 

The pure metal, however, is usually soluble in a dilute acid only 
with difficulty, unless a catalyst, such as a few drops of cupric 
sulphate, is added. The explanation of this phenomenon is 
probably electrical, since if a plate of pure zinc and a plate 
of copper are placed in dilute sulphuric acid and connected 
outside the solution by a wire, the zinc dissolves, hydrogen 
comes off from the copper, and an electric current flows 
through the wire. 

Zinc is used in the laboratory for the preparation of hydro- 
gen and also as a reducing agent. For the former purpose it 
is generally granulated by pouring the molten metal from a 
height into a tub of cold water, while for the latter purpose 
zinc dust is occasionally used. Zinc dust usually contains 
about 25 per cent, of zinc oxide, and if it is to be used quan- 
titatively must be estimated first. It is made by blowing a 
strong blast of air into melted zinc. 

Commercially, zinc is chiefly used to coat iron and thus 
prevent it from rusting. The iron object is first thoroughly 
cleaned and then dipped into molten zinc, when a thin layer 
of the latter metal is deposited as a coating on the iron. Iron 
treated in this way is called (inappropriately) galvanized iron. 
Zinc is also used in the preparation of alloys, e.g. brass and 
German silver (p. 227), and also in electric batteries. 


Zinc oxide may be made in any of the usual ways ; it is 
generally prepared by burning zinc in air. It is a white amor- 
phous powder, soluble in acids to form zinc salts, and turning 

ZINC 257 

yellow when heated, but regaining its white colour on cooling. 
It is used as a paint (" zinc white "), but has not the same 
covering-power as " white lead " (p. 328) ; it is not, however, 
blackened by sulphuretted hydrogen since zinc sulphide also 
is white. 

Zinc oxide is an amphoteric substance, since it dissolves in 
acids to form salts (basic property) and also in caustic alkalis 
(acidic property) ; the compounds formed in the latter case 
are called zincates and contain the zinc in the anion. 

ZnO + 2NaOH == Na 2 Zii0 2 + H 2 0. 

Zinc hydroxide is " insoluble " in water, but is not pre- 
cipitated by ammonium hydroxide if ammonium chloride ia 
present (Ionic Theory, p. 135). 

Zinc chloride, ZnCl 2 , is prepared anhydrous by passing 
hydrochloric acid gas over zinc heated in a tube. It cannot 
be prepared by heating the hydrated salt in air, as hydrolysis 
occurs resulting in the formation of zinc oxychloride 


ZnCl 2 + H 2 O = Zn< + HC1 * . 
X C1 

The " fused zinc chloride " that is commonly used as a 
dehydrating agent is made by dissolving zinc in hydrochloric 
acid and evaporating the solution, after concentration, in a 
stream of hydrochloric acid gas, to prevent hydrolysis. 

Zinc sulphate, ZnSO 4 , is usually met with in the hepta- 
hydrated form called white vitriol, ZnSO 4 .7H 2 O. Like all 
soluble zinc compounds, it is very poisonous. 

Zinc carbonate, ZnCO,, is made by adding a solution of 
sodium bicarbonate to a solution of a zinc salt ; sodium 
carbonate yields a basic zinc carbonate. 

Zinc sulphide, when not quite pure, is phosphorescent 
after exposure to light. It is largely used in work on X-rays 
and radio-activity, as when struck by a-particles, X-rays, etc., 
it is temporarily luminous (CROOKES, Spinthariscope). 

Zinc Alloys. For the composition of brass, see p. 227. 
High tensile brass is an alloy of copper, nickel and zinc. 



Naval brass, resistant to the corrosive action of sea-water, 
consists of Cu 63, Zn 36, Sn 1, parts by weight. There is 
reason to believe that copper and zinc form two definite 
compounds, viz. CuZn and CuZn 2 or Cu 2 Zn 8 . 

For certain important organic compounds of zinc, see text- 
books of organic chemistry. 


Group in Periodic System : II ; Symbol : Cd ; Valency : 
2; Atomic Weight: 11240; Melting Point : 321; Specific 
Gravity: 8-6. 

History. The name cadmia, according to DIOSCORIDES 
was given to one of the volatile products formed in the refine- 
ment of silver. Possibly owing to the mistake of a scribe in 
reading I for d (which letters are somewhat similar in Arabic), 
the name is met with in Arabian chemical treatises as calmia, 
whence our word calamine (ZnC0 3 ). Calmia or cadmia was 
probably either zinc oxide or sulphide ; the name was first 
used in its modern sense by STROHMEYEB (1817), who applied 
it to a metallic oxide that he found in the condensers of a 
zinc works. The metal was isolated shortly afterwards and 
called cadmium. 

Occurrence. Cadmium is almost always found in zinc 
ores, but occurs as sulphide (CdS) in a rare mineral called 

Extraction. When zinc ores are being smelted the cad- 
mium, being more volatile, comes over first and collects in the 
receivers as a brown powder, cadmium oxide, CdO. This is 
reduced with powdered coke. The mixture of oxide and coke 
is heated to not more than 800 C. in an iron retort, when 
cadmium vapour distils over and solidifies to a powder in a 
conical sheet-iron condenser. The metal is then purified by 
distillation or, more generally, by electrolysis. 

Properties. Cadmium is very similar to zinc ; its specific 
gravity is 8-6, melting-point 321, and boiling-point 770-780. 
Its soluble salts are only very slightly ionized. The oxide is 
a brown powder ; the sulphide, which comes down in Group 


II of the Analysis Tables, is bright yellow and is used as a 
paint ; it is insoluble in yellow ammonium sulphide. 

Cadmium is used in the Weston standard electric cell, and 
has recently found commercial application in cadmium - 
plating ; its importance for this purpose is likely to increase 
rapidly. It is also alloyed with silver in the jewelry trade, 
since the cadmium or silver alloy is more malleable and ductile 
than pure silver and also tarnishes less readily. 


Group in Periodic System : II ; Symbol : Hg ; Valency : 
1, 2 ; Atomicity of Vapour : 1 ; Atomic Weight : 200-6 ; 
Boiling Point : 357 ; Specific Gravity : 13-595. 

History. Mercury has been known from time immemorial ; 
the alchemists regarded it as the prime matter or " mother " 
of the other metals. PLINY distinguishes between the native 
metal, which he calls argentum vivum (" quicksilver ") and 
hydrargyrum (liquid silver) prepared from cinnabar (mercuric 
sulphide", HgS), by powdering the latter with vinegar in a 
copper vessel. Mercury was also prepared from cinnabar by 
strongly heating the ore in an iron vessel in a charcoal fire. 

The theory that metals are composed of sulphur and mer- 
cury was originated by the Arab chemists, and was very 
popular for many hundreds of years. In the symbolic lan- 
guage of the alchemists mercury was called by many names 
the poison, gambar, permanent water, the dragon, etc. 
" GEBER," the unknown author of many thirteenth century 
chemical treatises, describes it as follows 

" Argentvive, which aljo is called Mercury by the Ancients, 
is a vijcous Water in the Bowels of the Earth, by mojt tem- 
perate Heat united, in a total Union through its leajt parts, 
with the Jubjtance of white Jubtile Earth, until the Humid be 
contempered by the Dry, and the Dry by the Humid, equally. 
Therefore it eajily runs upon a plain Superficies, by reafon of 
its Watery Humidity ; but it adheres not, although it hath a 
vijcous Humidity, by reajon of the Drynefs of that which 
contemperates it, and permits it not to adhere. It is aljo (as 
Jome Jay) the Matter of Metals with Sulphur. And it eajily 


adheres to three Minerals, viz. to Saturn, and Jupiter, and Sol 
[that is, lead, tin, and gold], but to Luna [silver] more diffi- 
cultly. To Venus [copper] more difficultly than to Luna ; 
but to Mars [iron] in no wi/e, unlejs by Artifice. Therefore 
hence you may collect a very great Secret. For it is amicable, 
and pleajing to Metals, and the Medium of conjoyning Tinc- 
tures ; and nothing is Jubmerged in Argentvive, unlejs it be 
Sol. Yet Jupiter and Saturn, Luna and Venus, are dijjolved 
by it, and mixed ; and without it, none of the Metals can be 
gilded. It is fixed, and it is a Tincture of Eednefs of mojt 
exuberant Reflection, and fulgid Splendor ; and then it recedes 
not far from the Commixtion, until it is in its own Nature." 

Compounds of mercury, especially the oxide (HgO), sul- 
phide (HgS), and chloride (HgCl 2 , corrosive sublimate) have 
been known for at least 1,000 years. The name mercury was 
given to the metal on account of its fanciful connection with 
the messenger of the Gods, MERCURY or HERMES, who was 
also identified with the founder of chemistry, the Egyptian 
HERMES TRISMEGISTOS. From the earliest days the metallic 
nature of the element was recognized. 

Occurrence. Mercury is found in small quantities free or 
as an amalgam with silver. It is obtained chiefly from its ore 
cinnabar, HgS, which is mined in Spain, Mexico, Russia, 
Peru, China, Japan and other places. 

Extraction. Cinnabar (with or without small admixture 
of charcoal) is roasted in a current of air and the mercury 
vapour and sulphur dioxide produced are passed through 
cooled earthenware receivers called aludels (Arabic al-athdl). 
This method has been employed for 2,000 years, with various 
improvements in detail. 

If the cinnabar ore is rich, it is mixed with lime and the 
mixture strongly heated in cast-iron retorts, when the following 
reaction occurs : 

4HgS + 4CaO = 4Hg + CaS0 4 + 3CaS. 

Properties. Mercury is the only metallic element liquid 
at ordinary temperatures. It freezes at 39 C. and boils at 
357, giving a colourless vapour the molecules of which are 


monatomic (ratio of specific heats). The commercial metal 
is preserved in iron bottles (since it " adheres to Mars in no 
wise "), and usually contains a little copper and iron as 
impurities. Each bottle contains 75 Ib. of mercury. The 
mercury is purified by redistillation, and the traces of foreign 
metals still present are then removed by allowing the mercury 
to fall in tiny droplets through a long tube filled with dilute 
nitric acid, when the impurities are dissolved. If required 
of very high purity, the mercury is again redistilled in a 
vacuum or in carbon dioxide at low pressure. 

Mercury is not oxidized in the air at ordinary temperatures, 
but is gradually converted into the red oxide HgO, if heated 
in air or oxygen to a temperature just below its boiling-point. 
It is not dissolved by dilute sulphuric or hydrochloric acid ; 
concentrated sulphuric acid dissolves it, forming mercuric 
sulphate and sulphur dioxide 

Hg + 2H 2 S0 4 = HgS0 4 + 2H 2 + S0 2 , 

and so "does nitric acid, concentrated or dilute, forming mer- 
CUTOUS nitrate, HgNO 3 (cold dilute acid) or mercuric nitrate, 
Hg(N0 3 ) 2 (hot concentrated acid). Alkalis and most of the 
common gases (except chlorine) have no action on mercury. 

On account of its high density, low specific heat, low freez- 
ing-point and high boiling-point, it is extremely useful in 
thermometry ; the fact that it does not wet glass is an addi- 
tional advantage. It is also useful for the collection of gases 
that are soluble in water ; it was, indeed, by using mercury 
in the pneumatic trough that PKIESTLEY was able to collect 
ammonia. The important role played by mercury in the 
development of the oxygen theory of combustion will be at 
once remembered. 

Amalgams. Many metals will dissolve in mercury, and 
the solutions so formed are called amalgams (Arabic al-mula- 
gham, a mixture). Certain amalgams are used in the labora- 
tory, e.g. sodium amalgam, used as a reducing agent in 
presence of water. Sodium amalgam is made by gently 
warming mercury in an evaporating -dish and adding small 
pieces of sodium on the end of a knife. The sodium dissolves 


with evolution of heat and a flash of light. Dilute sodium 
amalgams are liquid, but more concentrated ones are crys- 
talline. Tin amalgam is used for making mirrors, while 
amalgams of cadmium, copper, gold, zinc, etc., are employed 
in dentistry for filling teeth, as they can be obtained as soft 
and plastic solids which gradually " set " or become hard. 
Mercury is used to form silver amalgam in the " amalgamation 
process " for the extraction of the latter metal (p. 229). One 
of the " artifices " referred to by " GEBER " for forming 
amalgam of " Mars " or iron is to grind iron filings with 
a paste of mercuric chloride and water. This process wai 
known to the Arab chemists of the twelfth century. 

COMPOUNDS OF MERCURY. Mercury in its compounds can 
be either univalent or bivalent ; it accordingly forms two 
series of compounds, the mercurous and mercuric. These may 
be regarded as derived from the two oxides, mercurous oxide, 
Hg 2 0, and mercuric oxide, HgO. 

Mercurous Compounds. 

Mercurous oxide, Hg 2 0, is a brownish -black powder made 
by the action of caustic soda solution upon mercurous 

Hg 2 Cl 2 + 2NaOH - Hg 2 + Hg 2 + 2NaCL 

It is unstable, and decomposes when warmed or when exposed 
to light, forming mercury and mercuric oxide. 

Mercurous chloride, Hg 2 Cl 2 , is a white powder often called 
calomel (Greek, beautiful black] because it is changed to a 
black substance (mixture of metallic mercury with NH 2 . Hg . Cl) 
by ammonia. Calomel is prepared by heating an intimate 
mixture of finely powdered mercuric chloride and mercury 
in an iron vessel provided with a lid ; the calomel sublimes and 
condenses as a white powder on the lid of the vessel. It still 
contains a little mercury and mercuric chloride as impurities, 
and is therefore washed first with dilute nitric acid, to remove 
the mercury, and then several times with distilled water, to 
remove the mercuric chloride, and the mercurous nitrate 
formed in the nitric acid washing. 

Mercurous chloride is also obtained, as a white precipi- 


tate, by adding a solution of hydrochloric acid or a chloride 
to a solution of a soluble mercurous salt, e.g. mercurous 

It is insoluble in water, and is used in medicine (as a pur- 
gative) and in the pottery trade. It dissociates on heating 
into mercury and mercuric chloride ; BAKER, however, has 
shown that perfectly dry calomel does not dissociate, and 
that the vapour density of the dry salt corresponds to the 
formula Hg 2 Cl 2 . Other work (cryoscopic and electrochemical) 
confirms this double formula. The alchemists prepared 
calomel mixed with corrosive sublimate by heating a mixture 
of mercury, salt, and vitriol. 

Mercurous nitrate (written Hg 2 (NO s ) 2 from analogy with 
the chloride) is obtained when mercury is dissolved in cold 
dilute nitric acid. It forms colourless crystals with two 
molecules of water of crystallization, Hg 2 (N0 3 ) 2 .2H 2 O, and 
is partially hydrolysed by water to a yellowish basic nitrate. 
The mercurous nitrate solution used as a reagent is therefore 
mixed with a little nitric acid, to prevent formation of the 
basic salt. 

Mercurous sulphate, Hg 2 S0 4 , is made by gently heating 
sulphuric acid with an excess of mercury. It is a practically 
insoluble crystalline powder, partially hydrolysed by water. 
It is chiefly of importance from its use in the standard electric 
cell (WESTON cell, see textbooks of physics). 

Mercurous iodide, Hg 2 I 2 , is a green powder made by rubbing 
iodine with excess of mercury, and one or two drops of alcohol, 
in a mortar. In the light it gradually decomposes into mer- 
cury and mercuric iodide, HgI 2 . 
Mercuric Compounds. 

Mercuric oxide, HgO, is obtained as a red crystalline powder 
by heating mercury in the air for several days to a tempera- 
ture just below its boiling-point. The earliest description 
of this preparation of mercuric oxide appears to be that given 
by an unknown Spanish Arab of about A.D. 1007, who in his 
book entitled RTTTBAT AL-HAKIM (" The Sage's Step "), says, 
" I took pure natural quivering quicksilver and placed it in 
a glass vessel shaped like an egg, and put this in another vessel 


shaped like a cooking-pot. I then set the latter on a very 
gentle fire and heated the apparatus for 40 days ; day and 
night. The original weight of the mercury was J Ib. At the 
end of 40 days I found it was all converted into a soft red 
powder, but I could detect no change in weight." It is 
formed as a yellow precipitate by addition of caustic soda to a 
solution of mercuric chloride 

HgCl 2 + 2NaOH = HgO + H 2 + 2NaCl. 

The difference in colour may be due to the difference in 
size of the particles of the oxide prepared in these two 
ways. Compare copper sulphate crystals and the powdered 
crystals. Alternatively it may be due to difference in crys- 
talline form. 

The red oxide may also be prepared by cautiously heating 
mercuric nitrate 

2Hg(N0 3 ) 2 = 2HgO + 4NO a + 0,. 

Strong heating of the nitrate results in metallic mercury being 
formed, from decomposition of the oxide. 

When mercuric oxide is heated, it turns black and splits up 
into oxygen and mercury (PRIESTLEY'S experiment on red 
calx of mercury). 

2HgO = 2Hg + 2 . 

It is soluble in dilute acids, giving mercuric salts. If it 
is boiled with a solution of ammonia it gives a yellow solu- 
tion of a compound called MILLON'S base, mercurammonium 
hydroxide, NHg 2 OH.H 2 O. This is used in testing for 

Mercuric chloride or corrosive sublimate, HgCl 2 , may be made 
synthetically by passing chlorine over heated mercury. This 
is the method employed commercially. It may also be made 
by heating a mixture of mercuric sulphate and common salt, 
or an intimate mixture of mercury, salt, ferric oxide, alum, 
and nitre, as used by the alchemists. Thus " GEBER " says : 
" Sublime Argentvive thus : R. of it lib. i. of Vitriol rubified 
[i.e. ferric oxide], lib. ij. Of Rock-Allom calcined, lib. i. of 
Common Sal, lib. ii. and of Salt-Peter one fourth part. Incor- 


porate all together, and Jublime : and gather the White, 
Denfe, and Ponderous, which Jhall be found about the Sides 
of the Veffel, and keep it, as we have appointed of other 
Things. But if, in the firjt Sublimation, you Jhall find it 
turbid, or unclean (which may happen, by reason of your own 
Negligence) again Jublime it with the Jame Feces, and rejerve 
it for vje." 

Mercuric chloride is a white crystalline solid, slightly 
soluble in cold water and more soluble in hot. It is used in 
medicine, as an antiseptic and germicide, in the form of a very 
dilute aqueous solution. It is sometimes employed as a 
perservative for the sleepers of railways, where they are liable 
to be attacked by insect pests. Like all soluble mercury 
compounds it is very poisonous. 

It is only slightly ionized in solution, and is a mild oxidizing 
agent ; thus, it oxidizes stamious chloride (SnCl 2 ) to stannic 
chloride (SnCl 4 ), sulphur dioxide in solution to sulphuric acid, 
and oxalic acid (in presence of a trace of iron salt as catalyst) 
to carbon dioxide, being itself reduced to mercurous chloride 
or even to mercury. 

Mercuric chloride reacts with ammonia under various con- 
ditions to give complex mercurammonium compounds, e.g. 

white precipitate, Hgx" , made by adding ammonia to 

NH a 
mercuric chloride solution. 

Mercuric nitrate [Hg(N0 3 ) 2 ] 2 .H 2 O or Hg(N0 3 ) 2 .pI 2 O, is 
obtained as colourless deliquescent crystals from the liquid 
produced by dissolving mercury in hot concentrated nitric 
acid. It is hydrolysed by water, giving a basic nitrate, 
Hg(N0 3 ) 2 .2HgO.H 2 O; boiling water carries the hydrolysis 
atill further, to mercuric oxide. 

Mercuric iodide, HgI 2 , is a characteristic red powder. It 
is made by adding potassium iodide to mercuric chloride 

HgCl 2 + 2KI = HgI 2 + 2KC1, 

and comes down first as a yellow precipitate ; the colour, 
however, rapidly changes to red. On heating to 130 the red 



form again changes to yellow ; the yellow is unstable at 
ordinary temperatures and slowly reverts to the red form 
rubbing or scratching accelerates the change. Mercuric 
iodide is soluble in excess of potassium iodide, forming the 
complex compound potassium mercuri-iodide, K 2 HgI 4 

2KI + Hgl, = K 2 HgI 4 . 

This shows none of the ionic reactions for mercury, since it 
ionizes into K", K', and HgI 4 ". An alkaline solution of 
potassium mercuri-iodide is called NESSLER'S solution, and is 
used as a test for ammonia or ammonium salts, with which 
it gives a yellow precipitate said to be mercurammonium 
iodide, Hg 2 N.I ; the composition of the precipitate is, how- 
ever, still uncertain. 

Mercuric cyanide, Hg(CN) 2 , may be made by dissolving 
mercuric oxide in dilute prussic acid and evaporating the 

HgO + 2HCN = Hg(CN) 2 + H 2 0. 

It forms large colourless crystals which do not ionize in 
solution and which decompose on heating, yielding cyanogen 
and mercury 

Hg(CN) 2 = Hg + C 2 N,. 

Mercuric sulphide, HgS, occurs naturally as the red mineral 
cinnabar. Formed by passing sulphuretted hydrogen through 
mercuric chloride solution it is black. The black form changes 
into the red on sublimation. It is made commercially (since 
it is used as a paint, vermilion) (i) by mixing mercury and 
sulphur with concentrated caustic potash solution and warm- 
ing to 50 ; (ii) by heating mercury and sulphur together, 
with constant and effective stirring. The crude sulphide is 
then purified by sublimation. 

Mercuric thiocyanate, Hg(CNS) 2 , is made by adding con- 
centrated potassium thiocyanate solution to mercuric chloride 
or nitrate solution. It forms a white precipitate, which when 
mixed with a little gum can be made up into the " eggs " of 
* PHARAOH'S serpents." When ignited these hatch out into 
awe-inspiring reptiles of a material which, very appropriately. 
is extremely poisonous. 



1. Compare and contrast the properties of magnesium and its com- 
pounds with those of (a) calcium and calcium compounds, (6) zino 
and zinc compounds. 

2. Describe the manufacture of magnesium. 

3. Give an account of the metallurgy of fcinc. 

4. What calcium compounds are used as fertilizers ? Write their 
formulae and describe their manufacture. 

5. Explain the use of strontium oxide in sugar -refining. 

0. Describe the preoaration and properties of the chlorides of 

7. Explain the importance of mercury in the chemical theories of 
the Middle Ages. 

8. What ere (a) the " eggs of Pharaoh's serpents," (b) Ndssler's 
solution, (c) cinnabar, (d) white precipitate ? Describe their prepara- 

9. What evidence is there for the constitution of bleaching-powder ? 
How would you estimate the available chlorine in this compound T 


TYPICAL ELEMENTS : Boron, Aluminium, 

Sub-group A : Scandium, Yttrium, Rare Eartli&. 
Sub-group B : Gallium, Indium, Thallium. 


Boron and aluminium resemble one another very little 
except in the fact that the formulae of their compounds are 
similar, since the valency of each element is 3. The element 
most closely related in chemical properties to boron is 
undoubtedly silicon. 

Boron is a typical non-metal, while aluminium is a true 

The " rare earth " metals are numerous and of com- 
paratively little importance chemically ; compounds of some 
of them are, however, used in the manufacture of incandescent 
gas-mantles and for other purposes. 

Scandium and gallium are of interest on account of 
MENDELEEFF'S prediction that they existed (p. 156) some 
years before they were actually discovered. 


Group in Periodic System : III ; Symbol : B ; Atomic 
Weight : 10-82 ; Valency : 3 ; Specific Gravity : 2-6-2-3 

History. The name buraq, whence our borax, was given 
by the Arabian chemists to many substances used as fluxes. 
One variety was tinkar or tinkal t obtained from Armenia and 



also from a salt lake in Tibet, and it is to this substance, 
Na 2 B 4 O 7 .10H 2 O, that the name " borax " is at present 
confined. In 1702 HOMBIIBG contributed a paper to the 
French Academy of Sciences, in which he described the pre- 
paration of a white flaky substance by adding sulphuric acid 
to a solution of borax. This new substance was called sal 
sedativum, and was largely used in medicine by the great 
French scientist LEMERY. Half a century later (1747), BARON 
showed that borax was a compound of sal sedativum and soda, 
and concluded that Homberg's substance was not a salt but 
an acid, for which the name boracic acid was suggested. On 
LAVOISIER'S oxygen theory of acids, boracic acid would be 
the oxide of a non- metallic element, and working from this 
point of view DAVY (1808), and GAY-LUSSAC and THENARD 
in the same year, succeeded in isolating the element boron, 
by heating boracic acid anhydride with potassium. 

Occurrence. Boron has not been found in nature in the 
free state % It occurs as crystalline boracic (or boric) acid in 
the mineral sassolino (Tuscany), but chiefly as borax or tinkal, 
Na 2 B 4 7 .10H 2 0, of which large deposits occur in California. 
Calcium borate, or colemanite, occurs in quantity in South 
America, while boric acid is found in the natural steam jets 
or soffioni of certain volcanic regions in Tuscany. 

Preparation. Boron proved difficult to isolate in the 
pure state. The simplest method of getting a reasonably 
pure specimen is to heat potassium borofluoride (KBF 4 ) with 
sodium in a stream of hydrogen 

KBP 4 + 3Na = KF + 3NaF + B. 

Reduction of the oxide (B 2 3 ) with magnesium powder has 
also been recommended (MOISSAN) 

B 2 3 + 3Mg = 3MgO + 2B. 

In each of these methods the product is treated with hydro- 
chloric acid, to remove the other substances, and the boron is 
left as a dark brown powder. 

Pure boron may be made, according to WEINTRAUB (1909) 
and PRING and FIELDING (1910), by striking the electric aro 


between copper electrodes in a mixture of boron trichloride 
vapour and excess of hydrogen 

2BC1 3 + 3H 2 ^= 2B + 6HC1. 

Most of the boron is deposited on the electrodes, and finally 
fuses into small beads, which fall off and solidify. 

WAKTH (1923) prepared boron of a high purity by reducing 
boron trichloride with hydrogen at 1,300-1,85Q in the pre- 
sence of a glowing filament of tungsten. 

Properties. Pure boron is an extremely hard black solid ; 
when broken it shows a conchoidal fracture (that is, the broken 
pieces show curved surfaces, and not the plane surfaces 
characteristic of crystalline bodies) and is therefore amorphous. 
Its specific gravity is 2-30-2-34 and its boiling-point about 
2,300. A remarkable property of boron is the rapidity with 
which its electrical resistance falls on rise of temperature. 
WEINTBAUB showed that a piece of boron which had a resist- 
ance of 5,620,000 ohms at 27 C. had a resistance of only 5 
ohms at a dull red heat. 

Pure boron does not oxidize in the air, even at very high 
temperatures ; it is, however, slowly oxidized to boric acid by 
concentrated nitric acid. It combines directly with nitrogen, 
if heated, to form boron nitride, BN, and with carbon (in the 
electric furnace) to form the carbide, B 4 C 3 . 

The brown powder obtained in the first two ways above 
described is an impure boron, which differs in many of its 
properties from pure boron. Colloidal boron has been pre- 
pared in several ways ; some of its colloidal solutions are brown 
and others are red. 

COMPOUNDS OF BORON. Boron is a non-metallic tervalent 
element ; it rarely exhibits a valency of 5. 

Boron hydrides. The simple hydride, BH 3 , is not known, 
but about ten others exist. The chief of these are B 2 H 6 , 
diboron hexahydride, and B 4 H 10 , tetraboron decahydride, a 
mixture of which is prepared by the action of hydrochloric 
acid on magnesium boride (made by heating boron trioxide, 
B 2 8 , with magnesium powder). B 2 H 6 is a colourless gas, 
while B 4 H 10 is a colourless volatile liquid with a very un- 

BORON 271 

pleasant smell. B 4 H 10 takes fire spontaneously on exposure 
to air or oxygen. 

Boron fluoride, BF 8 , is made by heating a mixture of 
calcium fluoride and boron trioxide with concentrated sul- 
phuric acid 

B 2 8 + 3CaF 2 + 3H 2 S0 4 = 2BF, + 3CaS0 4 + 3H 2 0. 

It is a colourless gas with a pungent smell, and must be 
collected over mercury, as it immediately attacks water, with 
formation of boric and hydrqfluoboric acids 

4BF 3 + 3H 2 = 3HBF 4 + H 3 B0 8 . 

The salts of hydrofluoboric acid are called borofluorides ; 
the most important is the potassium salt. 

Boron chloride, BC1 3 , is made by heating a mixture of 
boron trioxide and powdered charcoal in a stream of chlorine 

B 2 8 + 30 + 3C1 2 = 2BC1 3 + SCO, 

and passing the issuing gases through a U-tube surrounded 
by a freezing-mixture, when the boron trichloride condenses 
as a colourless fuming liquid boiling at 13. An alternative 
method (MAZZETTI and DE CARLI) is to heat iron boride 
(" ferroboron ") in a current of dry chlorine at 500. Water 
hydrolyses it irreversibly 

BC1 3 + 3H 2 O = H 3 B0 3 + 3HC1. 

Boron trioxide, B 2 3 (otherwise known as boric anhydride 
or boron sesquioxide), is obtained when boric acid is heated to 

2H 3 B0 8 = B 2 3 + 3H 2 0. 

It is a glassy colourless solid, very hygroscopic, and soluble 
in water, with which it combines to form boric acid. Al- 
though an acidic oxide, it shows some basic properties as well. 

Boric acids. Two boric acids are known ; the ordinary 
form is orthoboric acid, H 3 BO 3 , and this is converted into 
metaboric acid, HB0 2 , at a temperature of 100-140. Ortho- 
boric acid is volatile in steam, and, as previously mentioned, 
occurs in the soffioni or jets of steam that issue from the 
ground in certain parts of Italy. The steam is condensed in 


large tanks of water and the solution concentrated (largely 
by the heat of the natural steam) until the boric acid crystai- 
lizes out. The origin of the boric acid of the soffioni is 
uncertain ; it has been suggested that the steam may traverse 
a stratum of rock containing boron sulphide, boron nitride, or 
tourmaline (a boron mineral), all of which yield boric acid in 
contact with steam. 

In the laboratory, boric acid is conveniently made by 
addition of sulphuric acid to a hot concentrated solution of 
borax ; on cooling, the boric acid separates as a pearly crystal- 
line solid, greasy to the touch. It is a very weak acid, so weak 
that it has no action on methyl orange ; it may, however, be 
titrated with caustic soda if excess of glycerol is added to the 
solution and phenolphthalein is used as indicator. Under 
these conditions it behaves as a monobasic acid 

H 3 B0 8 + NaOH = NaB0 2 + 2H 2 0. 

Boric acid or boracic acid finds many uses in everyday life 
it is a well-known antiseptic, and was formerly used as a 
preservative for cream, etc., although it is harmful if taken 
internally. Its use as a preservative has now been forbidden 
by Act of Parliament. 

Borax, or sodium pyroborate, Na 2 B 4 7 .10H 2 0, occurs 
naturally as tinkar or tinkal, but is also made from the South 
American colemanite, Ca 2 B O n .5H 2 0, by boiling it for some 
time with sodium carbonate solution 

Ca 2 B 6 O n + 2Na 2 C0 8 - 2CaC0 3 + Na 2 B 4 7 + 2NaB0 2 . 
Carbon dioxide is then blown through the solution, when 
4NaB0 2 + C0 2 = Na 2 C0 3 + Na 2 B 4 7 . 

In this way all the colemanite is converted into borax, which 
is crystallized out from the filtered solution by evaporation. 

Borax is sodium pyroborate decahydrate, Na 2 B 4 7 . 10H 2 (X 
On heating it swells up and loses its water, forming anhydrous 
borax, which on melting goes to a clear transparent glass. 
This glassy substance readily dissolves metallic oxides, in 
nianv cases with production of characteristic colours. 
(" Borax Bead " test.) 


Borax is still used as a flux, as it has been for the last 1,000 

In solution, borax is largely hydrolysed, 

Na 2 B 4 O 7 + 7H 2 O ^= 2NaOH + 4H 3 B0 8 . 

The solution therefore has an alkaline reaction, and since 
methyl orange is not affected by the weak boric acid, a solution 
of borax may be titrated with an acid as if it were caustic soda 
solution, if methyl orange is used as indicator. 

The green flame produced when boric acid (or a borate), 
sulphuric acid, and alcohol are heated together and the alcohol 
vapour ignited, is due to the formation of an organic boron 
compound, ethyl borate, (C 2 H 6 ) 3 B0 8 , which burns with the 
characteristic greenish flame. 


Group in Periodic System : III ; Symbol : Al ; Atomic 
Weight: 27-0; Valency: 3; Melting-point: 659; Specific 
Gravity : 2-7. 

History. The word alum is derived from the Latin 
alumen, " a mineral salt with an astringent taste." Alum 
occurs in many parts of the ancient world, and was known 
to the Arabian chemists, who usually placed it in the group 
of compounds called by the general name of buraq or borax. 
Other early chemists more correctly classed it with the 
vitriols with which it has obvious chemical similarities, 
although it is not isomorphous with them. PARACELSUS, 
however, was unable to extract any metal from alum and 
therefore stated, in his usual dogmatic fashion, that it was 
not a vitriol. No further elucidation of the nature of alum 
was made until 1746, when POTT concluded that it was 
a compound of a new " earth." Eight years later 
MARGGRAF prepared this earth, " alumina," from clay, and 
showed that it was quite distinct in properties from lime. 

DAVY attempted to isolate a metal from alumina, which 
he quite rightly assumed to be a metallic oxide ; he was 
unsuccessful, however, and it was not until 1827 that metallic 
aluminium was prepared. In that year W6HLER heated 



anhydrous aluminium chloride with potassium and obtained 
aluminium as a white metal of low specific gravity. 

Occurrence. Metallic aluminium is not found in nature, 
but aluminium compounds are extremely abundant and are 
found universally in rocks and clays. A few of the chief 
naturally occurring compounds of aluminium are the fol- 

Silicates, e.g., clay, slate, mica, felspar (KAlSi 3 8 ), kaolin 

(Al 2 Si 2 7 2H 2 0), garnet, topaz, tourmaline. 
Oxides, e.g. bauxite (A1 2 3 .2H 2 0), corundum (A1 2 8 ), 
diaspore (A1 2 3 .H 2 0), spinel (MgAl 2 4 ), chrysoberyl 
(BeAl 2 4 ). 
Fluoride, e.g. cryolite (Na 3 AlF 6 ) ; phosphate, e.g. turquoise. 

Fro. 57. Cross Section of Hall Cell for Manufacture of Aluminium. 

A. Fused electrolyte. B. Layer of powdered coke. C. Molten aluminium. D. Carbon 
lining of cell. E. Plug for running off the fused metal. 

Manufacture. Aluminium was first manufactured by 
heating sodium aluminium chloride, NaAlCl 4 , with sodium, 
a process devised by DEVILLE in 1854 

NaAlCl 4 + 3Na = 4NaCl + Al. 

Shortly afterwards the fluoride, cryolite, Na 3 AlE 6 , was used 
instead of the chloride, but the high price (60 per Ib.) of the 
aluminium made in this way rendered the commercial use of 
the metal impossible. 

In 1886, however, an electrolytic process was invented 
simultaneously by HALL in America and by H^ROULT in 


France. These chemists discovered that alumina (A1 2 3 ) 
would dissolve in fused cryolite and that the solution so 
obtained could be electrolysed, oxygen coming off from the 
anode and aluminium being liberated at the cathode. In 
effect, therefore, the process is an electrolysis of alumina, since 
the cryolite is left unchanged. 

Alumina is obtained from the mineral bauxite (A1 2 3 . 2H 2 0), 
which is found in Ireland, France and the United States. 
Natural bauxite always contains iron oxide and silica as 
impurities, and since the presence of iron and silicon in 
aluminium is very harmful (rendering it easily attacked by 
water, etc.), the bauxite has to be carefully purified before 
use. Realization of this fact, followed by preparation of a 
purer aluminium, at once led to a considerable improvement 
in the financial conditions of aluminium companies. 

Purification of the bauxite is carried out in many different 
ways. In one of these the crushed bauxite is dissolved in 
hot concentrated caustic soda solution, forming sodium 

A1 2 O 3 .2H 2 O + 2NaOH = 2NaA10 a + 3H 2 O. 

The impurities are insoluble and are removed by passing the 
liquid through a filter-press. To the filtrate a little freshly 
prepared gelatinous aluminium hydroxide is added, when 
practically all the sodium aluminate is gradually decomposed 
and the aluminium precipitated as hydroxide. This is filtered 
off and converted into alumina by ignition, while the filtrate 
of caustic soda solution is used again. 

The cryolite is obtained from the large deposits that occur 
in Greenland. It is fused in an iron box lined with gas carbon, 
which forms the cathode. The anode consists of a number of 
stout carbon rods which dip into the molten cryolite ; the 
latter forms a shallow layer which in the Hall process is only 
some 6 inches in depth. Alumina, purified as already 
described, is dissolved in the fused cryolite, fresh supplies 
being added from time to time as electrolysis proceeds. 
Molten aluminium collects on the floor of the cell and is run 
off as required. The carbon anodes are burnt to carbon 


monoxide and dioxide by the oxygen liberated % and have to 
be replaced ; as they are expensive, attempts have been made 
to use other substances for the anodes, but hitherto without 
success. Analysis of the anode gases indicates that the 
formation of the aluminium and oxidation of the anode 
proceed roughly according to the equations 

A1 2 3 + 3C - SCO + 2A1. 

Properties. Aluminium is a metal of a bluish white 
colour. It can be highly polished and is malleable and ductile, 
although it becomes brittle and can be powdered at a tem- 
perature just below its melting-point. It has a low specific 
gravity, and as it is fairly hard, is extensively used in airship 
production and for other purposes where " lightness " is a 
prime requisite. It is also employed on a large scale for the 
manufacture of domestic utensils such as saucepans. First- 
grade commercial aluminium is of about 99 per cent, purity ; 
the chief impurities are iron and silicon. 

Aluminium melts at 658 and boils at 1,800 ; its specific 
gravity is 2-7. It is stable in the air, apparently owing to the 
formation of a thin film of oxide which protects the under- 
lying metal from further oxidation. On heating aluminium 
in the air it burns with a brilliant white flame, forming 
alumina. Nitric acid has no action upon it ; it dissolves, 
however, in hot concentrated sulphuric acid, forming alu- 
minium sulphate and sulphur dioxide 

2A1 + 6H 2 S0 4 - A1 2 (SO 4 ) 8 + 6H 2 O + 3S0 2 . 

It dissolves easily in hydrochloric acid, forming the chloride 
and hydrogen 

2A1 + 6HC1 = 2A1C1 3 + 3H 2 . 

It also dissolves very quickly in solutions of caustic alkalis, 
to form aluminates, with evolution of hydrogen 

2H 2 + 2A1 + 2NaOH = 2NaA10 2 + 3H f . 
Sodium al animate. 

Aluminium is a powerful reducing agent ; in the form of 


powder it readily reduces metallic oxides with evolution of 
intense heat. A mixture of ferric oxide and aluminium 
powder is called thermite ; when ignited by means of a piece 
of burning magnesium ribbon the whole mass becomes incan- 
descent and molten iron is left. This reaction is employed 
for welding two pieces of iron together, e.g. tram-rails, 
without the necessity of removing them from their positions. 
Thermite was used during the war for filling incendiary bombs, 
and a mixture of aluminium powder with ammonium nitrate 
formed the Mills' bomb explosive ammonal. 

Many alloys of aluminium are of importance. The chief are 
duralumin (94 per cent. Al, 0-5 per cent. Mg, 4-5 per cent. Cu, 
0-75 per cent. Mn, 0-25 per cent. Fe), magnalium (95 per cent. 
Al, 4-5 per cent. Mg, 0-5 per cent. Sn, etc.), aluminium bronze 
(90 per cent. Cu, 10 per cent. Al), and aluminium brass (30 per 
cent. Zn, 69 per cent. Cu, 1 per cent. Al). 

Aluminium amalgam is readily attacked by water, yielding 
mercury, hydrated alumina or aluminium hydroxide, and 
hydrogen. The aluminium-mercury couple (Al/Hg) consists 
of aluminium foil covered with a film of aluminium amalgam. 
With methyl alcohol it forms a reducing agent widely used in 
organic chemistry, e.g. in the preparation of methane from 
methyl iodide. 


Aluminium sesquioxide, or alumina, A1 2 3 , occurs 
naturally in the anhydrous state as corundum, and as hydrates 
in diaspore, A1 2 O 3 .H 2 0, bauxite, A1 2 3 .2H 2 0, and hydrar- 
gillite, A1 2 3 .3H 2 O. Coloured varieties of corundum are 
known as ruby (red), oriental topaz (yellow), oriental emerald 
(green), oriental sapphire (blue), and amethyst (violet). The 
colour is due to the presence of small quantities of foreign 
metallic oxides such as those of chromium, manganese, cobalt, 
etc. The preparation of artificial rubies and sapphires from 
aluminium oxide is now an important industry, millions of 
carats being produced every year. To make rubies, fused 
alumina is mixed with a little chromium sesquioxide, Cr 2 3 ; 
for the production of sapphires a mixture of magnetic oxide of 


iron and titanium dioxide is used instead of th$ chromium 

Impure corundum is called emery and is used in making 
" emery-paper." 

Pure aluminium oxide may be made by heating the pure 
hydroxide or nitrate. Alumina is usually soluble in dilute 
acids, but if it is strongly heated it appears to change into 
another modification at 850, and this is insoluble. 

Aluminium hydroxide, A1(OH) 3 , is formed as a white 
gelatinous precipitate on adding caustic alkali or ammonia 
to a solution of an aluminium salt ; it is insoluble in excess 
of ammonia but dissolves in excess of soda or potash to form 
aluminates. As it is also soluble in dilute acids, giving 
aluminium salts, it is an example of the class of amphoteric 
substances ; in solution (it is very slightly soluble) it ionizes 
in two ways 

(i) A1(OH) 3 ;= AT" + 30H'. 
(ii) A1(OH) 3 ;== 3H* + A1(Y". 
It is used in the dye industry as a mordant. 
The aluminates of sodium and potassium, made by dis- 
solving aluminium or aluminium hydroxide in caustic soda 
or potash solution, are often represented by the formulae 
Na 3 A10 3 and K 3 A1O 3 . Cryoscopic determinations, however, 
show that in solution they probably have the formulae 
NaA10 2 and KA10 2 , and ought therefore more properly to 
be called meto-aluminates. They have been prepared in the 
solid crystalline state and the crystals have the following 
compositions: potassium salt, KA10 2 .1JH 2 0, sodium salt, 
NaA10 2 .2H 2 0. 

Several anhydrous crystalline meta- aluminates occur 
naturally : they are called spinels 
magnesia spinel, Mg(A10 2 ) 2 . zinc spinel, Zn(A10 2 ) 2 . 
iron spinel, Fe(A10 2 ) 2 . 

Minerals of similar constitution, in which the aluminium 
has been replaced by tervalent iron, chromium, or magnesium, 
are also called spinels. 
Aluminium chloride, AlCl t , is made anhydrous by heatr 


ing aluminium in a current of dry chlorine or hydrochloric 
acid gas 

2A1 + 3C1 2 - 2A1C1 3 ; 

or by heating an intimate mixture of alumina and carbon in 

It is a white deliquescent solid, often tinged with yellow by 
ferric chloride present as impurity. When heated, it sublimes 
at 180 ; the molecules of the vapour up to 400 are A1 2 C1 6 , 
but at this temperature rapid dissociation begins and at 700 
is complete 

A1 2 C1 6 ;== 2A1C1,. 

In solution in certain organic solvents such as ether the 
molecules are A1C1 3 . 

Aluminium chloride is very soluble in water, which partially 
hydrolyses it 

A1C1 3 + 3H 2 O ;= A1(OH) 3 + 3HC1. 

Crystals of aluminium chloride hexahydrate, A1C1 3 .6H 2 O, 
can be obtained by cautiously evaporating a solution of the 
metal, oxide, or hydroxide in hydrochloric acid, or by saturat- 
ing an aqueous solution of the chloride with hydrochloric acid 
gas. On heating, the hexahydrate loses water and hydro- 
chloric acid and is converted into alumina. 

The anhydrous salt is commonly used as a catalyst in 
certain organic reactions (e.g. Friedel and Crafts' reaction). 

Aluminium sulphate, A1 2 (S0 4 ) 3 can be prepared as a 
hydrate with 18H 2 by evaporating to crystallization a solu- 
tion of aluminium hydroxide in dilute sulphuric acid 

2A1(OH) 8 + 3H 2 30 4 - A1 2 (S0 4 ) 3 + 6H 2 O. 

The crystals of the hydrate lose their water on heating, and 
the anhydrous sulphate is left. 

Aluminium sulphate is of importance commercially since 
it is employed in the purification of sewage and also in paper- 
making. It is manufactured by heating china-clay (kaolin) 
or bauxite with sulphuric acid and is put on the market in 
solid blocks called " alum-cake." 

With the sulphates of the alkali-metals, aluminium sul- 


phate forms a series of double salts called the alums ; they 
have the general formula R 2 S0 4 . A1 2 (S0 4 ) 3 .24!! 2 0, where 
R = (Li), Na, K, Rb, or Cs. The existence of lithium alum 
is doubtful. Ordinary " alum " is potassium alum. 

Similar compounds are known in which the aluminium 
sulphate is replaced by the sulphate of other tervalent metals 
such as iron and chromium. Potassium chrome alum, for 
instance, is 

K 2 S0 4 .Cr 2 (S0 4 ) 8 .24H 2 0, 

and ammonium ferric alum 

(NH 4 ) 2 S0 4 . Fe(S0 4 ) , . 24H 2 0. 
The general formula for all " alums " is therefore 

R 2 S0 4 .R' 2 (S0 4 ) 3 .24H 2 0, 

where R a univalent metal, usually an alkali-metal, and 
R' a tervalent metal, usually iron, chromium or 

Potassium alum is made from alum-cake by dissolving it 
in water, adding the appropriate weight of potassium sul- 
phate, and crystallizing the solution. It is also made from 
a mineral called alunite, K 2 S0 4 .A1 2 (S0 4 ) 3 .4A1(OH) 3 , found 
in France, Hungary and other countries, by treatment with 
dilute sulphuric acid. Sufficient potassium sulphate is added 
to convert the excess of aluminium sulphate thus formed into 
the double salt, and the solution then evaporated. 

Another source of alum is alum-shale, a slate (or aluminium 
silicate) containing iron pyrites. The shale is roasted for 
about a week and then treated in the same way as alunite. 

Clay, Porcelain and China. An account of the clays 
and their uses in the pottery trade falls outside the scope of 
this elementary book. The student who is interested in the 
subject should read the appropriate articles in THOBPE'S 
Dictionary of Applied Chemistry, etc. 

Ultramarine. The soft blue mineral known as lapis- 
lazuli is a complex sodium aluminium silicate containing about 
13 per cent, of sulphur. It is now manufactured, under the 
name of " ultramarine," by heating a mixture of china-clay, 


charcoal, soda-ash and sulphur. It is used in washing clothes 
(" blue-bag "), to make them look white. 


1. Give an account of the history of boron. 

2. Compare and contrast the properties of boron with those of 

3. Write a description of boric acid and borax. How is borax made 
commercially ? 

4. Mention the chief naturally occurring aluminium compounds. 

5. Describe the manufacture of aluminium. 

6. Boron and aluminium are classed together in the Periodic Sys- 
tem. How far is this classification supported by the chemical relation- 
ships of the two elements ? 

7. When sodium sulphide is added to a solution of alum, aluminium 
hydroxide is precipitated. Can you explain this ? 

8. Describe the preparation and properties of anhydrous aluminium 


TYPICAL ELEMENTS, Carbon, Silicon. 

Sub-group A : Titanium, Zirconium, Cerium, Thorium. 
Sub-group B : Germanium, Tin, Lead. 


Although lead and tin show certain resemblances to one 
another, they are very different from the characteristic non- 
metals carbon and silicon. Carbon and silicon, again, have 
the same valency, 4 (and therefore form compounds of similar 
formulae), but otherwise differ markedly from one another. 
In chemical properties silicon is most closely related to 

All the elements in Group IV are quadrivalent, and many of 
them may be bivalent as well. 

Chemically, cerium should be classed with the rare -earth 
elements of Group III. 


Group in Periodic System : IV ; Symbol : C ; Valewcy : 
4; Atomic Weight: 12-00. 

History. The history of carbon was fairly uneventful 
until the beginning of the nineteenth century, when, with the 
aid of the Atomic Theory, the constitution of " organic " or 
carbon compounds was first successfully studied. Since that 
time the chemistry of the compounds of carbon has become 
an entire subject in itself and is called Organic Chemistry. 

In the form of coal, which crops out at the surface of the 



earth in many places, carbon has been known for at least 
2,000 years. A Roman altar, excavated at Bath, was found 
to have coal ashes upon it, and there are many mentions of 
coal in the ancient writers. It was not until 1775 that the 
diamond was proved to be a form of carbon ; in that year, 
however, LAVOISIER, and twenty-two years later TBNKANT, 
showed that when the diamond is burnt carbon dioxide is 
produced, the latter proving in addition that the weight of 
carbon dioxide formed from a given weight of diamond is the 
same as that which would be given by the combustion of an 
equal weight of charcoal. 

In 1800 MACKENZIE proved that graphite, formerly confused 
with molybdenum sulphide under the name of " blacklead," 
was another crystalline form of carbon. 

Occurrence. Carbon is found free in nature in three 
allotropic forms diamond (crystalline), graphite (crystalline), 
and coal (impure amorphous form). In combination it is 
found in metallic carbonates (limestone, marble, dolomite, 
etc.) ; as carbon dioxide, in the air ; as hydrocarbons (com- 
pounds of carbon and hydrogen only) in petroleum and 
natural gas ; and in the form of more complex bodies in all 
living tissue, of which, indeed, it is the prime constituent. 

Properties. The three allotropic forms will be con- 
sidered in turn. 

(i) Diamond 

Diamonds are found in many parts of the world. In olden 
days they came chiefly from India, but most of them are now 
obtained from South Africa, Brazil, and Australia, although 
a few are still extracted in India and Borneo. In South 
Africa diamonds are found occasionally in alluvial river sand, 
but chiefly in " pipes " or inverted cones of a blue clay lying 
vertically in the earth and supposed to have been formed by 
volcanic activity in a remote geological age. To mine this 
clay, a shaft is sunk through the rock at one side of the 
" pipe," and horizontal tunnels are then made from the shaft 
into the " earth." The " earth " is brought to the surface 
and exposed to the weather for several months, during which 


it falls to a coarse powder. This powder is then agitated in 
a stream of water, which removes the lighter particles, and 
the concentrated earth is then washed over a layer of grease, 
to which the diamonds adhere. 

Natural diamonds differ considerably in size, colour, 
and value. The largest diamond ever discovered was the 
Cullinan, which was found in the Transvaal in 1905 and before 
cutting weighed 3,025f carats over a pound and a quarter (1 
carat = 3-17 grains or 0-2054 gm.). Colourless diamonds are 
the most valuable ; the " black diamonds " (carbonado) are 
used for making glass-cutters, etc., and for cutting and 
polishing the colourless stones. 

For purposes of jewelry, the natural diamonds are cut in 
duch a way that as much internal reflection as possible is 
caused ; this process requires great skill, and many valuable 
stones have been spoilt through inexpert cutting. There is 
always bound to be some loss in weight on cutting a stone, of 
course, but in some cases it is found necessary to cut the* stones 
down severely. Thus the Koh-i-noor (" mountain of light ") 
which originally weighed 186 carats, had to be cut down to 
106 carats. 

The density of the diamond is 3-5 and its refractive index 
2-45 ; it is the hardest substance known. It is extremely 
stable towards chemical reagents ; acids have no effect upon 
it, but it is slowly attacked by fused sodium carbonate, 
forming carbon monoxide 

Na 2 C0 3 + C = Na 2 O + 2CO. 

When a diamond is heated in air to about 800 it takes fire 
and burns brilliantly, forming carbon dioxide. If heated in 
the electric furnace in absence of air it swells up, and changes, 
partially at least, into graphite. 

It is generally believed, from a study of the rocks anc? <jlay 
in which they occur, that diamonds have been formed by the 
crystallization of carbon from a solution of that element in 
iron, or perhaps in a silicate, possibly under great pressure. 
This conclusion is supported by the chemical researches of the 
versatile French chemist MOISSAN. Many chemists had 


previously been attracted to the fascinating problem of the 
artificial production of diamonds, but success was first 
attained by Moissan in 1893. He took a small tube of soft 
iron fitted with a screw cap, filled the tube with pure sugar- 
carbon (p. 294), and screwed down the cap so that the carbon 
inside was strongly compressed. He dropped this tube 
into a crucible of molten iron heated in the electric furnace ; 
the crucible and contents were then removed and quickly 
cooled by immersion in melted lead. Under these conditions 
the iron on the outside of the mass solidified while that inside 
was still liquid. Now iron, like water, expands on solidifica- 
tion, so that when the interior of the mass solidified it did so 
under the extremely high pressure caused by the central iron 
trying to expand inside the hard solid crust. 

When cold, the iron was dissolved away in hydrochloric 
acid, and a number of crystals of carbon were left. Some of 
these proved to be graphite, but others were diamonds, both 
black and colourless. The diamonds obtained in this way are, 
however, very small (the largest was scarcely more than half 
a millimetre in diameter), and the process is therefore not a 
commercial success, as the cost of making the diamonds is 
about three times their market value when made. 

The experiments of SIR CHARLES PARSONS (1907, 1918, etc.) 
appear to show that retention of carbon monoxide, and not 
production of high pressure, is the essential part played by the 
outer crust of iron in Moissan's process ; experiments made to 
melt carbon under enormous pressures in the electric furnace 
always gave negative results. 

(ii) Graphite 

The name " graphite " was given to this substance on 
account of its use as a writing material (y()d<f)(O, I write). It 
was called plumbago or black lead possibly because, like lead, 
it marks paper, or because it was supposed to be a lead com- 
pound. Graphite occurs in large quantities in many different 
localities : Cumberland, Bohemia, Ceylon, Siberia and Cali- 
fornia are the chief. 

It is manufactured on a large scale at Niagara by a process 



invented by AOHBSON (who also discovered carborundum, 
p. 313). About three tons of a mixture of powdered coke 
with a little sand and some pitch are placed in a brick furnace 
the bottom of which is protected by a layer of sand. Through 
the walls of the furnace project two stout carbon rods, con- 
nected to a dynamo. On passing a powerful current through 
the mass for 24 to 30 hours a very high temperature is main- 
tained and the coke is converted into a pure and soft graphite, 
called Acheson's graphite. This can be finely powdered and 
is then used in the preparation of aquadag and oildag. 

FIG. 58. Furnace for the Manufacture of Graphite. 

(an aqueous solution of deflocculated ^4cheson 
graphite) is a colloidal suspension (p. 145) of graphite in a dilute 
solution of tannin. The tannin is a protective and prevents the 
precipitation of the graphite. Aquadag is used as a lubricant. 

Oildag is made by filtering aquadag through rubber (ordinary 
filters are useless for filtering colloids) and mixing the paste of 
colloidal graphite obtained in this way with lubricating oil. 
This yields a colloidal solution of graphite in oil, called oildag, 
extensively . used as a lubricant. 

Finely powdered graphite itself may also be used as a 
lubricant ; many people " blacklead " their bicycle chains 
instead of putting oil on theiw, while if the vulcanite mouth- 


piece of a pipe fits very stiffly into the wooden stem a smoker 
will rub his " lead " pencil on the part of the vulcanite which 
fits into the wood, and so ease it. 

Graphite crystallizes in lustrous black hexagonal plates 
that feel greasy to the touch ; it is a good conductor of heat 
and electricity, and is very stable to chemical reagents. A 
mixture of nitric acid and potassium chlorate will act slowly 
upon graphite (BRODIE, 1855), to form a green substance 
which on treatment with acidified potassium permanganate is 
converted into an amorphous yellow compound called graph- 
itic acid. Graphitic acid is usually given the formula C^H^Os 
but it is really a complex mixture. Nitric acid has no action 
on diamonds, and converts amorphous carbon into mellitic 

The specific gravity of graphite is variable ; it is usually 
about 2-3. Graphite finds many applications. Powdered 
graphite mixed with clay is used for making pencils. A 
mixture of sand and graphite is made up into refractory 
crucibles capable of withstanding a high temperature and the 
action of many acids. 

Graphite is also used to protect iron and steel (e.g. fenders, 
stoves, and grates) from rust, and as a polish for gunpowder. 
On account of its comparatively good conductivity it is often 
used in electrolytic processes. 

(iii) Amorphous Carbon 

Coal. Coal has been formed by the decomposition of 
plant remains by bacteria and saprophytic fungi, in absence 
of air and under high pressures. Recent experiments (BER- 
GIUS, 1913-1925) have shown that a substance very similar 
to coal may be made by heating peat, cellulose and water 
to 300-400 under a pressure of 5,000 atmospheres ; it is 
therefore possible that heat changes in the earth's crust may 
have had something to do with coal-formation. 

Various stages in the conversion of vegetable matter into 
coal are shown by the substances called peat, lignite, bitumin- 
ous coal, cannel coal (or " parrot " coal), and anthracite. 
Ordinary coal is the bituminous variety. 



The approximate composition of these substances is given 
in the following table 


Per Cent. 

Per Cent. 

Per Cent. 

Per Cent. 










Bituminous coal .... 
Cannel coal . . 





Anthracite . * 





The figures in the table are rough only, and variation within 
fairly wide limits is found in practice. The table serves, 
however, to show two important facts 

(i) The percentage of carbon from peat to anthracite 
steadily increases. 

(ii) The percentage of hydrogen falls off much less rapidly 
than that of the nitrogen and oxygen (the figures given in the 
last column but one represent chiefly oxygen ; the percentage 
of nitrogen remains fairly constant at about 1 to 2). 

Peat is a brown substance composed of the partially decom- 
posed remains of bog plants. It is used as a fuel and is cut in 
many parts of the country, as on Sedgmoor for instance. It 
represents in all probability the first stage in the formation of 

Lignite, or brown coal is generally considered to represent 
the next stage. It is harder than peat and usually contains 
a slightly higher percentage of carbon. It differs from peat 
and coal in that it is mainly composed of fossil wood \ how- 
ever, there seems to be no doubt that in the course of ages it 
would have become converted into a coal. 

Bituminous coal is composed of the remains of cryptogamic 
or flowerless plants (thus resembling peat, chiefly composed 
of decayed mosses) which flourished in an amazing manner 
in the Carboniferous Age. These flowerless plants are now 
represented on the earth only by insignificant descendants 
such as the ferns, horse-tails and club-mosses. 


The problem of the chemical constitution of coal is now 
engaging the attention of chemists ; it is too complicated a 
subject to be discussed here. Coal certainly contains free 
amorphous carbon, but it contains in addition many complex 
carbon compounds. According to STOPES, four distinct 
constituents may be detected in coal by examination of thin 
sections under the microscope ; these are called durain, 
clarain, fusain and vitrain. In general, coal appears to consist 
of a colloidal mass in which are embedded (i) woody substances, 
(ii) the decayed spores of flowerless plants, (iii) resinous matter 
and (iv) a heterogeneous debris formed of i, ii and iii. Further 
light has been thrown on the subject by a chemical examina- 
tion of the substances dissolved out of coal by heating it with 
benzene under pressure. 

Cannel coal is so called because a piece of it if lit will burn 
like a candle. It is chiefly used for making coal-gas. 

Anthracite is a hard coal which may contain as much as 97 
per cent, of carbon. It is used for heating boilers (hence the 
name steam-coal) and also in anthracite stoves. Anthracite 
is difficult to ignite, but when ignited goes on burning very 
steadily, without flame, and with production of intense heat ; 
anthracite stoves are therefore economical in use and require 
very little attention. Much of the world's anthracite is 
mined in South Wales ; the Admiralty pits produce anthra- 
cite for the Navy, the ships of which are now, however, largely 
oil- driven. 

Distillation of Coal. The destructive distillation of coal 
is more conveniently dealt with in the study of Organic 
Chemistry ; the student is therefore referred to the author's 
Outlines of Organic Chemistry (London, Edward Arnold 
& Co.). 

When coal is heated in iron retorts with exclusion of air, 
four main products are formed 

(i) Coke (left in the retorts). 

(ii) An aqueous liquid from which the ammonium sulphate 
of commerce is made ; the " ammoniacal liquor" 

(iii) Coal-tar. 

(iv) Coal-gas. 



(i) Coke. This consists chiefly of carbon (about 70-80 per 
cent.) ; it also retains some of the nitrogen of the original 
coal, the rest appearing in the form of the ammonia of the 
ammoniacal liquor. Coke is used as a fuel and as a source 
of gaseous fuels (water-gas and producer-gas, pp. 301-3) ; its 
use in metallurgical operations is referred to many times 
in the course of this book. 

(ii) Ammoniacal liquor. For the treatment of this 
liquid see p. 339. 

(iii) Coal-tar. Tar is a black liquid containing many 
important substances such as benzene, toluene, naphthalene 
and phenol. It is redistilled, and the distillate, which is 
collected in fractions, is worked up in various ways that are 
described in the books on organic chemistry. 

(iv) Coal-gas. The crude gas which passes on from the 
receivers in which the ammoniacal liquor and tar are con- 
densed still contains traces of these and other substances, and 
has to undergo purification. It is first passed through a 
number of vertical pipes, the condensers, in which most 
of the remaining tar and ammoniacal liquor are condensed 
and collect at the bottom. It is then usually sent through a 
tar-extractor, in which any tar that may have escaped condens- 
ation is removed. 

From the extractor the gas passes into the exhauster, where 
the pressure of the gas is adjusted by means of pumps, and 
thence into the scrubbers, which are iron towers packed with 
broken bricks or coke down through which water trickles. 
In the scrubbers the last of the ammonia is removed, and 
some of the sulphuretted hydrogen and carbon dioxide is 

The gas now passes through the purifiers, which are iron 
chambers containing trays filled with moist ferric oxide (" bog 
ore," from Holland and Belgium). Here the sulphuretted 
hydrogen is removed, since it reacts with the ferric oxide to 
form ferric sulphide, ferrous sulphide, and sulphur 

2Fe a O t + 2H 2 + 6H 2 S = Fe a S 8 + 2FeS + S + 8H 2 0. 

When the ferric oxide has all been changed according to the 



above equation, the product is called " spent oxide." This ia 
taken out and exposed to the air, water being added if neces- 
sary. Oxidation with separation of sulphur occurs and the 
spent oxide is now said to be " revivified " 

2Fe 2 S 3 + 30 2 - 2Fe 2 3 + 68. 
4FeS +30 2 = 

The mixture of iron oxide and sulphur is returned to the 
purifiers and used repeatedly until it contains about 50-60 per 
cent, of sulphur ; it is then sold to a lead-chamber sulphuric 
acid works and used as a source of sulphur dioxide. 

Carbon disulphide, which is always present as an impurity 
in coal-gas, is sometimes removed and sometimes allowed to 
remain. It can be absorbed by using slaked lime instead of 
ferric oxide in the purifiers, since it combines with the calcium 
hydrosulphide formed by the action of sulphuretted hydrogen 
on the lime 

Ca(OH) 2 + 2H 2 S - Ca(HS) 2 + 2H 2 0. 
Ca(HS) 2 + CS 2 = CaCS 3 + H 2 S. 

The product, calcium thiocar donate, CaCS 3 , has no commer- 
cial application and is difficult to deal with. It has been 
shown, however, that by passing coal-gas over finely divided 
nickel at 250-300 reduction of the carbon disulphide by the 
hydrogen present is brought about 

CS 2 + 2H 2 = 2H 2 S + C. 

This is another interesting application " of SABATIER and 
SENDERENS' reducing agent (reduced nickel and hydrogen : 
the nickel is a catalyst). 

From the purifiers the gas passes into the gasometers. It 
consists chiefly of hydrogen, methane and carbon monoxide, by 
the combustion of which most of the heat of a coal-gas flame 
is produced, together with small quantities of acetylene, 
benzene and ethylene, to which the luminosity of the flame is 
due. There are also present varying amounts of gaseous 
impurities. The following table (BUTTERFIELD, 1913) gives 
tho analysis of typical specimens of coal-gas 







Carbon monoxide . 


Unsaturated hydrocarbons (acetylene, etc.) 

2-5- 5-0 

0-0- 3-0 

0-0- 1-5 

Of recent years it has become the practice to add a certain 
amount of water-gas (p. 301) to the pure coal-gas. 

Gas -carbon is a very hard form of carbon deposited as a 
lining on the retorts in gas-works. It is formed by the 
decomposition of hydrocarbons by heat, and as it is a good 
conductor of electricity is used for making the " carbons " 
for arc -lamps. 

Soot and lampblack are also forms of amorphous carbon. 
Lampblack contains oil and other substances as impurities. 
It is used in making printer's ink. 

Animal charcoal or bone-black is a mixture of calcium 
phosphate and about 10 per cent, amorphous carbon made by 
the destructive distillation of bones. It is used in sugar 
refining, since when boiled with a coloured solution it will 
often absorb the colouring matter and leave the solution 
colourless ; the brown solution of crude sugar is decolourized 
in this way. 

Wood charcoal is a porous form of amorphous carbon made 
by burning wood with insufficient air (" charfeoal-burning "). 
It is important in the laboratory chiefly on account of its 
remarkable power of absorbing gases ; thus 1 c.c. of charcoal 
will absorb nearly 200 c.c. of ammonia at ordinary temperature 
and pressure, and almost as much sulphur dioxide. 

Moreover, SIR JAMES DEWAR discovered that this absorp- 
tive power of charcoal was greatly increased at low tempera- 
tures, a fact which has found several useful applications, e.g. 



production of high vacua, separation and purification of the 

rare gases. 
Pure carbon is prepared by heating pure sugar in air until 

it is completely charred and reheating the charred mass first 

in a current of chlorine and then in a current of hydrogen. 

The product is called sugar-carbon and is extremely pure. 
COMPOUNDS OF CARBON. Over 1,000,000 compounds of 
carbon have been described, and 
there seems to be no limit to the 
number which might be prepared. 
The study of these compounds is 
the province of organic chemistry ; 
we shall describe here only a few 
of the most important. 

HYDROCARBONS are compounds 
consisting of hydrogen and carbon 
only. Three of the simplest are 
methane, ethylene and acetylene. 

Methane, or marsh -gas, CH 4 , 
is produced by bacterial action 
upon vegetable matter at the bottom 
of ponds, marshes, etc., hence one 

C- \^ _f\ * * ts names ' I* a l so occurs in coal- 
\ mines and is there called fire- 
'-"" damp. In the laboratory it is often 

of prepared (in an impure state) by 
heating anhydrous sodium acetate 
with sodalime (p. 246). 

CH,.COONa + NaOH = Na a C0 3 + CH 4 . 
A purer gas may be made by adding water to aluminium 

A1 4 C 8 + 12H a O = 4A1(OH) 8 + 3CH 4 . 
Another convenient method of obtaining the pure gas is 
to dissolve dry magnesium in a dry ethereal solution of methyl 
iodide, forming methyl magnesium iodide (CH 8 .Mg.I), and 
then to decompose the latter substance with dilute hydro- 
chloric acid : 

J. 60. Absorption 
Gases by Carbon. 



2HC1 = 2CH 4 


2CH 8 .Mg.I + 2HC1 = 2CH 4 + Mgl a + MgCl t . 

Methane is a colourless gas, with no taste or smell ; it In 
insoluble in water and burns with a practically non-luminous 

CH 4 + 20 2 = C0 2 + 2H 2 0. 

A mixture of methane with air or oxygen in certain proportions 
will explode if ignited. 

A mixture of 1 volume of methane with 1 volume of chlorine 
may explode if exposed to bright sunlight, the products 

Fia. 61. Preparation of Methane. 


being methyl chloride (a gas) and hydrochloric acid, 
reaction takes place in ordinary daylight. 

CH 4 + C1 2 = CH 8 C1 + HC1. 

Under suitable conditions, this substitution process may be 
continued until all the hydrogen of the methane has been 
replaced by chlorine 

CH 3 C1 + C1 2 = HC1 + CH 2 C1 2 , dichloromethane. 

CH 2 Cl a + C1 2 = HC1 + CHC1 8 , trichloromethane or chloro- 

CHOI, + Cl t == HC1 + CC1 4 , tetrachloromethane or carbon 




Ethylene, C 2 H 4 , is obtained as a colourless gas when ethyl 
alcohol, C 2 H 5 OH, is heated with strong sulphuric acid ; 
sulphovinic acid, C 2 H 6 HSO 4 , is first formed but immediately 
splits up, on heating, into ethylene and sulphuric acid 
C 2 H 5 OH + H 2 S0 4 = C 2 H 5 HS0 4 + H 2 0. 
C 2 H 6 HS0 4 - C 2 H 4 + H 2 S0 4 . 

It will be noted that, theoretically, there is no loss of 
sulphuric acid in the reaction, which resolves itself into the 

FIG. 62. Preparation of Ethylene. 

decomposition of the alcohol by the acid into ethylene, 
C 2 H 4 , and water 

C 2 H 6 OH = C 2 H 4 + H 2 0. 

[You should be careful to avoid the inexcusable blunder of 
saying that the acid dehydrates the alcohol, i.e. takes water 
out of it. Alcohol does not contain water, and therefore water 
cannot be removed from it. It contains the elements of water, 
but that is a different thing. A factory does not contain a 
house, yet if you take the factory to pieces you may be able to 
build up a house from certain of them. Remember that in 
science looseness of expression is a serious fault.] 


Ethylene is more conveniently made by passing the vapour 
of ethyl alcohol over aluminium oxide heated to 360 ; the 
alumina acts as a catalyst 

C 2 H 5 .OH = C 2 H 4 + H 2 0. 

Ethylene is a colourless gas with a sweetish taste and smell. 
It burns with a luminous flame and yields an explosive mix- 
ture with air. One of its most characteristic properties is 
that it will combine directly with bromine or chlorine to form 
colourless oils, ethylene dibromide and ethylene dichloride. 

C 2 H 4 + Br 2 = C 2 H 4 Br 2 ; C 2 H 4 + C1 2 = C 2 H 4 C1 2 . 

These substances are additive compounds, and should be 
contrasted with the substitution compounds formed by 
methane. Substances that will form addition compounds 
are called unsaturated (e.g. ammonia, which " adds on " 
hydrochloric acid NH 3 + HC1 = NH 4 C1). 

Acetylene, C 2 H 2 , is well-known as the gas made by the 
action of water on calcium carbide 

CaC 2 + 2H 2 = Ca(OH) 2 + C 2 H 2 . 

As usually prepared it has an objectionable smell, due to the 
presence of traces of phosphine (from the calcium phosphide 
in the carbon) ; the. pure gas has a sweetish smell, not unlike 
that of ethylene. It may be synthesized by passing an 
electric current between two carbon electrodes in an atmo- 
sphere of hydrogen (Fig. 63). 

Acetylene is an unsaturated compound more so than 
ethylene, since 1 molecule of it will combine directly with 
2 molecules of bromine or chlorine 

C 2 H 2 + 2C1 2 - C 2 H 2 C1 4 . 

It burns with a smoky flame unless burnt in special burners, 
designed to supply plenty of air to the flame ; from these 
burners it is delivered either as a fine jet or as a thin sheet and 
then burns with an intensely hot and brilliant flame. The 
flame of acetylene burning in oxygen is as hot as the electrio 
arc (3,500) and is used for welding and steel-cutting (oxy- 
acetylene flame). Acetylene is of growing commercial import- 




ance as the starting-point in the synthetic production of many 
organic substances. 

A mixture of acetylene and oxygen explodes with great 
violence when ignited. Acetylene is an endothermic com- 

When acetylene is passed through ammoniacal cuprous 
chloride it gives a chocolate brown precipitate of cuprous 
acetylide, Cu 2 C 2 ; ammoniacal silver nitrate with acetylene 
gives the corresponding silver acetylide, Ag 2 C 2 . Both these 
compounds are explosive, especially the silver one. 


' '.Hydrogen and 

FIG. 63. Synthesis of Acetylene. 

Oxides of Carbon. Carbon forms several oxides, of which 
the chief are carbon sub-oxide, C 8 O 2 , carbon monoxide CO, 
and carbon dioxide, CO 2 . 

Carbon suboxide, C 3 2 , is a colourless gas made by the 
action of phosphorus pentoxide upon an organic compound 
called ethyl malonate 
/COOC 2 H 6 

CH 2 + 2P 2 5 = C 3 2 + 2C 2 H 4 + 4HP0 8 . 

\COOC 2 H 5 
It will burn in air, forming carbon dioxide, and dissolves in 


water to form malonic acid, of which it is therefore an anhy- 


C S 2 + 2H 2 = CH a 


Carbon monoxide, CO, may be made in many ways, such 
as the following 

(i) Carbon dioxide is passed over red-hot coke (Fig. 64) 

C0 2 + C = 2CO. 
(ii) Steam is passed over red-hot coke 

C + H 2 = CO + H a . 

The mixture of carbon monoxide and hydrogen obtained in 
this way is called water -gas (p. 301). 

Dilute solution of 
sodium carbonate 

FIG. 64. 

(iii) Formic acid is heated with concentrated sulphuric 

H.COOH = H 2 + CO. 

Sodium formate may be used instead of formic acid. 
(iv) Oxalic acid is heated with concentrated sulphuric acid 
and the evolved gases passed through a Drechsel bottle con- 
taining caustic soda solution to absorb the carbon dioxide 
that is also formed 

| = CO + CO, + H,0. 


(v) Potassium ferrocyanide is heated with concentrated 
sulphuric acid (care ! dilute acid gives prussic acid, HCN). 

K 4 Fe(CN) 6 + 6H 2 + 6H 2 S0 4 

= 2K 2 S0 4 + 3(NH 4 ) 2 SO i + FeS0 4 + 6CO. 

Methods iii and iv are most convenient for laboratory pur- 
poses. Carbon monoxide is very often formed when a metallic 
oxide is reduced with carbon, e.g. 

ZnO + C = Zn + CO. 

Carbon monoxide is a colourless gas with a practically 
unnoticeable smell ; as it is extremely poisonous and con- 
stantly present in coal-gas its lack of smell makes it a source 
of considerable danger. Many cases of carbon monoxide 
poisoning are reported in the newspapers. The poisonous 
action of the gas is due to the fact that carbon monoxide 
combines with the red colouring-matter (hcemoglobin) of 
the blood, forming a bright red compound called carboxy 
hcemoglobin. "The haemoglobin is therefore rendered incapable 
of carrying on its function as a distributor of oxygen to all 
parts of the body, and death results from a kind of suffoca- 

Carbon monoxide is a colourless gas which can be liquefied 
by cold and pressure to a colourless liquid boiling at 190 
and freezing to white crystals at 200. It will burn in the 
air with a blue lambent flame, forming carbon dioxide. A 
mixture of 2 volumes of carbon monoxide and 1 volume of 
oxygen will explode if sparked, giving 2 volumes of carbon 
dioxide. Hence, by Avogadro's Hypothesis, 2 molecules of 
carbon monoxide + 1 molecule of oxygen give 2 molecules of 
carbon dioxide. But the molecule of carbon dioxide is C0 
and that of oxygen is 2 , 

.'. molecule of carbon monoxide must be CO. 

Carbon monoxide is an unsaturated compound, since the 
carbon atom in its molecule is only bivalent whereas the 
normal valency of carbon is four. Carbon monoxide will 
therefore readily combine with chlorine or oxygen and is thus 
a powerful reducing agent (cf. manufacture of iron, p. 545). 


If passed into an ammoniacal solution of cuprous chloride, 
carbon monoxide is absorbed, with formation of the com- 
pound CuCl. CO .2H 2 ; this has been isolated in the form of 
unstable white crystals. 

A characteristic property of carbon monoxide is its power 
of forming addition compounds, carbonyls, with gently heated 
metals. Certain of these, such as nickel carbonyl, Ni(CO) 4 
(p. 562), are of considerable importance. They are generally 
liquids which split up again into metal and carbon monoxide 
at higher temperatures. 

When carbon monoxide is passed over solid caustic soda 
heated to 200 the two combine to form sodium formate (cf. 
method iii, p. 299). 


This is a commercial method of obtaining sodium formate 
and thence formic acid itself. 

The heat of formation of carbon monoxide is 29,000 calories, 
and its heat of combustion 67,960 calories. It is clear that 
carbon monoxide will therefore be a very efficient gaseous fuel 
and as such is very widely used. 


Producer Gas is a mixture of carbon monoxide (30 per cent.) 
and nitrogen (62 per cent.) made by sending a blast of air 
through white-hot coke. If used directly for combustion, 
both the heat of formation and heat of combustion of the 
gas are available, but if the gas is allowed to cool first the heat 
of formation is naturally dissipated and only the heat of 
combustion will then be available ; there will thus be a loss 

29 000 

of heat in the proportion of , or about 30 per cent. 

F F 96,960 * 

Water Gas is the name given to a mixture of gases, chiefly 
hydrogen (49 per cent.), carbon monoxide (42 per cent.) and 
carbon dioxide (4 per cent.), made by passing steam over 
white-hot coke 

C + H 2 - CO + H 2 . 

This action is strongly endothermic (29,000 calories) ; the 


coke therefore cools down rapidly. After the steam-blast has 
been on for 5 or 6 minutes it is cut off and an ait-blast turned 
on, when producer-gas is formed with evolution of heat and 
the coke becomes incandescent again. The water-gas and 
producer-gas are usually stored in separate gasometers. 

In an alternative process sufficient air is used in the air- 
blast to form carbon dioxide ; this is of course useless, but 
the loss in coke is more than counterbalanced by the extra 
evolution of heat 

C + oxygen > CO + 29,000 calories. 

C + O a -> C0 a + 96,960 calories. 

Semi-water Gas is a mixture containing hydrogen (17 per 
cent.), carbon monoxide (27 psr cent.), and nitrogen (52 per 
cent.), made by passing air and steam together over white-hot 
coke. The heat evolved in the reaction is sufficient to keep 
the coke at a suitable temperature. 

Carburetted Water Gas is made by spraying paraffin into the 
steam-blast used in the water-gas process. It burns with a 
luminous flame, whereas water-gas burns with a blue non- 
luminous flame. 

Coal-gas is often mixed with water-gas at the present day ; 
the practice has the disadvantages, from the consumer's point 
of view, that the gas is thereby made much more dangerously 
poisonous, and that the heat produced by combustion of the 
gas is much less than that given out by the combustion of an 
equal volume of pure coal-gas. The latter difficulty has been 
overcome by selling gas on the basis of its heating-power (i.e., 
at so much per therm) instead of by volume. 

Water-gas has recently been used for the commercial pre- 
paration of hydrogen. A mixture of water-gas and steam is 
passed over reduced nickel, iron or copper at a temperature 
of 400-500, sometimes under slightly increased pressure ; the 
following reaction then occurs, the metal acting as a catalyst 

CO + H 2 = C0 2 + H a . 

The hydrogen is freed from the carbon dioxide by absorb- 
ing the latter in lime. Water-gas is also largely employed in 
the manufacture of methanol (methyl alcohol, CH 3 OH). It 


t* mixed with half its volume of hydrogen and the mixture is 
fieated to 450, under a pressure of 200 atmospheres, in the 
presence of certain metallic oxides that act as catalysts : 

CO + 2H f = CH 8 OH. 

[For a useful account of the manufacture of coal-gas, water- 
gas and carburetted water-gas, see H. HOLLINOS, The Manu- 
facture of Gas, School Science Review, September and Decem* 
ber, 1927.] 

Carbon dioxide, C0 2 . Carbon dioxide was discovered by 
VAN HELMONT (1577-1644), who called it gas sylvestre (" the 
wild gas of the woods "). It is interesting to note that the 
word gas was coined by van Helmont and is defined by one of 
his English disciples, GEORGE THOMSON (Direct Method of 
Curing Chymically, London, 1675), as " a wild invisible spirit, 
not to be imprisoned or pent up, without damage of what 
conteins it, arising from the Fermentation of the concourse of 
some Bodies, as it were eructating or rasping this untamable 
Matter." Two more of Thomson's definitions seem to be 
worth quoting 

" Chymist is one who imitates Nature in Separating the 
Pure Juice from the Dross and Filth for the use of Medicine 
Mechanicks, and the advancement of Mettals." " Chymico- 
phant, one who seems to be a Chymist, but is not really." 

Carbon dioxide was rediscovered in 1754 by JOSEPH BLACK, 
who called it fixed air and showed that it was liberated from 
chalk by the action of a dilute acid. Its synthesis from carbon 
and oxygen, and therefore its composition, were first shown 
by LAVOISIER in 1781 . Lavoisier re-named it acide carbonique t 
although it is, of course, not an acid but the anhydride of an acid. 

Carbon dioxide is constantly present in the atmosphere, of 
which it forms about 0*03 per cent, by volume. It is also 
found in solution in all natural waters, and is found in large 
quantities in ravines in certain volcanic districts, e.g. the 
Grotto del cane at Naples and the Valley of Poison in Java. It 
is a product of combustion of coal and other carbonaceous 
fuels, and also of the respiration of plants and animals. 

Preparation. Although carbon dioxide is formed when 



carbon is burnt in air or oxygen, it is usually prepared in tha 
laboratory by the action of dilute hydrochloric acid upon 
marble or other carbonates 

CaC0 3 + 2HC1 = CaCl 2 + H 2 + CO 2 . 
It may be dried by sulphuric acid and collected by down- 
ward displacement, or over mercury. A convenient appa- 
ratus for this purpose is that 
known as Kipp's Apparatus (Fig. 

It is formed by strongly heating 
the carbonates of heavy metals, 

FIQ. 65. 

and also during the fermentation 
of sugar solution by means of 

i. C 12 H 22 O n + H 2 = C 6 H 12 0, 
+ C 6 H 12 6 . 

ii. C 6 H 12 6 = 2C0 2 + 2C 2 H 6 OH. 

Alcohol remains in solution. 
Properties. Carbon dioxide is a 
colourless gas usually said to be 
odourless ; it has, however, a dis- 
tinct but faint smell. It will not 
burn and will not allow things to 
burn in it. Although it is not 
poisonous it will not suppdrt life ; 
the unpleasant effects of "stuffy " 
rooms, however, are not due to 
the accumulation of carbon di- 
oxide in them, but rather to the 
stagnation of the atmosphere. Even a comparatively high 
concentration of carbon dioxide causes little discomfort if 
the air is kept in motion by a fan. 

Carbon dioxide can easily be liquefied. The liquid is 
colourless and under atmospheric pressure at once solidifies 
fco a white crystalline mass of solid carbon dioxide. At 


atmospheric pressure, then, carbon dioxide has no meltjng-point 
or boiling-point ; rise of temperature merely causes sublima- 
tion of the solid, and not fusion. A mixture of solid carbon 
dioxide and ether is a good freezing-mixture, giving a tempera- 
ture of 79. If carbon dioxide is allowed to escape, from 
a cylinder of the compressed gas, into a flannel bag, the bag 
quickly fills with solid carbon dioxide, the so-called " carbon 
dioxide snow." This is used in ice-cream factories and carts. 
Carbon dioxide is soluble in water, one volume of water 
dissolving an equal volume of the gas at 15. The solution is 
weakly acid and contains a little carbonic acid, H 2 C0 3 , which 
has never yet been isolated 

Carbonic acid is an extremely weak acid ; thus an aqueous 
solution of carbon dioxide turns litmus only to a " port -wine " 
colour and has no effect upon methyl orange. It will, how- 
ever, decolourize pink phenolphthalein. Carbonic acid is a 
dibasic acid and "forms both normal carbonates and bicar- 
bonates. When carbon dioxide is passed into a solution of 
an alkali it forms first the normal carbonate and then the 

(i) 2NaOH + C0 2 = Na 2 C0 3 + H 2 0. 
(ii) Na 2 C0 8 + H 2 + C0 2 = 2NaHC0 3 . 

Sodium bicarbonate is neutral in solution, since it ionizes 
into Na" and HC0 3 ', and the further dissociation of HCO/ 
into H* and C0 3 " is negligibly small. Sodium carbonate, on 
the other hand, is alkaline in solution, owing to hydrolysis 

(i) H 2 ;= H- + OH'. 
(ii) Na 2 C0 3 ^r Na' + Na + C0 3 ". 
(iii) C0 3 " + IT ^= HCO/, thus leaving free hydroxyl 
ions, which render the solution alkaline. 
Or, Na 2 C0 3 + H 2 ;== NaOH + NaHCO 3 . 

All bicarbonates are soluble in water and yield the norma) 
carbonates on heating 

2NaHC0 8 - Na 2 C0 3 + H 2 + C0 a . 


When carbon dioxide is passed into lime-water it gives first 
a milky precipitate of calcium carbonate 

Ca(OH) 2 + C0 a = CaC0 3 + H 2 0. 

Further action of carbon dioxide will turn the milky liquid 
clear again owing to the formation of soluble calcium bicar- 

CaC0 3 + H 2 + C0 a = 

The presence of calcium bicarbonate in natural water makes 
the water " temporarily hard " (p. 446). 

Dissociation of Calcium Carbonate. When calcium car- 
bonate is heated it splits up into calcium oxide and carbon 
dioxide. If the reaction is carried out in a closed vessel (so 
that the carbon dioxide cannot escape), connected with a 
pressure-gauge or manometer, it is found that a* state of 
equilibrium is set up, which may be represented by the 

CaC0 3 ^= CaO + C0 a . 

In terms of the phase rule (p. 146), we have 3 phases 
(CaO, CaC0 8 , C0 2 ), and two components (CaO, C0 2 ). 

P + F = C + 2, or 3 + F = 2 + 2. 

.*. F = 1, or the system is uni variant. Therefore at any 
given temperature there will be a definite dissociation- 
pressure of carbon dioxide. 

In terms of Le Chatelier's principle, since an increase ol 
volume takes place on dissociation we should expect increase 
of pressure to diminish the dissociation ; this is found to be 
the case. 

Percarbonates. Electrolysis of a cooled saturated solution 
of potassium carbonate yields a bluish- white, amorphous, 
deliquescent precipitate of potassium percarbonate, K 2 C 2 O e . 
This is a strong oxidizing agent. Other percarbonates are 

Composition of Carbon Dioxide. If a weighed quantity of 
carbon is burnt in oxygen and the weight of carbon dioxide 



produced is estimated, by absorbing the gas in a weighed bulb 
containing caustic potash solution and finding the increase in 
weight, the gravimetric composition of carbon dioxide can be 
calculated. Experiments have shown that the ratio of carbon 
to oxygen in carbon dioxide is exactly 
3:8. The vapour density of the gas 
is 22, .*. its molecular weight is 44. The 
atomic weight of carbon is 12 and 
that of oxygen is 16, .'. the formula 
of the gas must be C0 2 . 

Volumetrically the composition of 
carbon dioxide may be determined by 
burning a piece of carbon in a measured 
volume of oxygen (Fig. 66). After 
cooling it will be found that there 
is no change in volume ; hence 1 
volume of oxygen is contained in 1 
volume of carbon dioxide. 

.'. by Avogadro's Hypothesis, 1 
molecule of carbon dioxide contains 
1 molecule of oxygen, 

.'. formula is C^Oj. 

x is found from the vapour density 
of the gas and the atomic weight of 

Carbon dioxide, in the Air. See p, 

Carbonyl chloride, COC1 2 , or 
phosgene, is made by passing a mix- 
ture of equal volumes of carbon mon- 
oxide and chlorine over gently heated 
animal charcoal 

CO + Cl a = COC1,. 

It is a colourless and very poisonous gas which reacts 
with water to form carbon dioxide and hydrochloric 

COC1, + H 2 = C0 a + 2HC1. 

Fio. 66. 


It can be regarded as the acid chloride of carbonic acid 

/OH /Cl 


\OH \C1 

Carbonic acid. Carbonyl chloride. 

'Carbon disulphide, CS 2 , is formed by passing sulphur 
vapour over red-hot coke, or by heating sulphur and coke 
together in an electric furnace 

C + 2S - CS 2 . 

It is a colourless, volatile, and very inflammable liquid 
which has a pleasant sweetish odour when pure. It generally 
contains impurities, however, which give it a very evil smell. 
It boils at 46 and melts at 110. 

Carbon disulphide vapour is poisonous and has been used 
as an insecticide and rat-killer, but its inflammable nature 
makes its use a matter of no little risk. It is a very good 
solvent for rubber and other organic substances, but has been 
displaced very largely of late years by the non-inflammable 
acetylene tetrachloride, C 2 H 2 C1 4 , and carbon tetrachloride, 
CC1 4 . 


Group in Periodic System : IV ; Symbol : Si ; Atomic 
Weight : 28-3 ; Valency : 4 ; Specific Gravity : 2-35-2-5. 

History. Silica (Si0 2 ) was regarded by BECKER (1655) 
as an " earth J> [i.e. a basic substance, like lime], which he 
called terra vitrescibilis owing to its use in glass making. 
TACHENTUS, however, in 1660, showed that it would dissolve 
in fused caustic alkalis, and therefore must possess an acidic 
nature. LAVOISIER regarded silica as the oxide of a new 
element, a hypothesis which was supported by SCHEELE'S 
work on silicon fluoride (SiF 4 ) and the silicic acids. GAY- 
LUSSAC and TH^NARD, working on Lavoisier's theory and 
Scheele's facts, in 1811 passed the gaseous silicon fluoride 
over heated potassium and succeeded in preparing elementary 


Occurrence. Silicon is very widely distributed, and 
forms more of the earth's crust than any other element except 
oxygen. It is always found as the oxide, silica, Si0 2 . This 
may occur in the free state as quartz (" rock-crystal ") (12 per 
cent, of the earth's crust) or as silicates, which are compounds 
of silica with basic metallic oxides. The most widely- distri- 
buted silicate is that of aluminium, which is the chief con- 
stituent of many rocks and also of clay. Many specimens of 
quartz show delightful' colours and are prized as gems, e.g. 
jasper (red), carnelian (red), amethyst, chrysoprase (green), 
cairngorm (yellow). The opal consists of amorphous silica, 
while agate is a variety of quartz. Flints are composed of a 
mixture of quartz and amorphous silica (hence th& name 
silica, from silex, a flint). Sand is a form of silica. 

Preparation. Silicon may be obtained in the amorphous 
state by heating potassium silicofluoride, K 2 SiF 6 , with 

K 2 SiF 6 + 4K - 6KF + Si (BERZELIUS, 1823), 
or by heating finely-powdered sand (SiO 2 ) with magnesium 

Si0 2 + 2Mg = Si + 2MgO. 

The crystalline variety may be made by melting the 
amorphous form and allowing the liquid to crystallize, or 
directly by passing silicon tetrachloride vapour over fused 
aluminium and dissolving out the residual excess of aluminium 
in hydrochloric acid 

3SiCl 4 + 4A1 = 3Si + 4A1C1 3 . 

Other methods have been described, of which perhaps the 
most important is that of heating sand and carbon in the 
electric furnace. If too much carbon is present, silicon 
carbide or carborundum is formed. 

Properties. Amorphous silicon is a brown powder, which 
melts at about 1,550 and can be boiled in the electric furnace. 
It can be made to burn in the air if strongly heated, and 
burns readily in oxygen, forming silica, Si0 2 . It combines 
with many elements at a high temperature, forming silicides. 
The specific gravity of amorphous silicon is 2-35. 


It is insoluble in acids (except hydrofluoric mixed with a 
little nitric), but dissolves in fused or aqueous caustic alkalis 
to form a silicate and hydrogen 

Si + 2NaOH + H 2 - Na 2 Si0 8 + 2H 2 . 

Crystalline silicon is a yellow or brown solid ; it is not so 
reactive as the amorphous variety. 

All specimens of silicon hitherto prepared have been im- 
pure, so that its true properties are not known with accuracy 
or certainty. 


Silicon dioxide, Si0 2 , is the most important silicon com- 
pound. It occurs naturally in the crystalline form as rock- 
crystal ; other varieties have been mentioned previously. 
Silica is also found in plants and animals the sharp cutting 
edge of blades of grass is often strengthened with silica, while 
skeletons of tiny plants called diatoms form large deposits of the 
" siliceous earth " known as kieselguhr, used in the manufacture 
of dynamite, etc. (p. 360). 

Silica occurs in three crystalline forms and also in the 
amorphous state. The crystalline varieties are known as 
quartz, tridymite, and cristobalite. When an acid is added to 
a solution of sodium silicate, a white gelatinous precipitate of 
a silicic acid is obtained ; if this precipitate is heated it loses 
water and yields amorphous silica. 

Silica is an acidic oxide, and in combination with water 
forms the silicic acids. Of late years it has become very im- 
portant in the manufacture of " quartz glass " ; when silica 
is heated it becomes soft and can be made into tubes, basins, 
crucibles, etc., and since silica has a very low coefficient of 
expansion, quartz-glass vessels can withstand great and sudden 
changes of temperature without cracking. A silica basin, for 
example, may be made red hot and cooled under the tap. 
Fused silica may also be drawn out into very fine threads used 
in mirror-galvanometers, etc. 

Vitreosil is an opalescent form of quartz glass, made by 
passing a powerful electric current through a rod or plate of 


carbon packed in sand. It is cheaper than the transparent 
form but equally good for most purposes, although of course 
it does not look so attractive. 

Silicic acids. It is supposed that metasilicic acid, H 2 SiO,, 
and orthosilicic acid, H 4 SiO 4 , exist, but the existence of neither 
has been definitely proved. When hydrochloric acid is added 
to an aqueous solution of sodium metasilicate, Na 2 Si0 8 , a white 
gelatinous precipitate is obtained, which after drying in a 
desiccator has a composition approximating to that of meta- 
silicic acid, H 2 Si0 8 . A substance which has roughly the 
composition H 4 Si0 4 or Si(OH) 4 , orthosilicic acid, is formed, 
according to NORTON and ROTH (1879), by hydrolysing silicon 
chloride, SiCl 4 , with water and washing the precipitate with 
benzene and dry ether. 

The silicates are more important than the silicic acids. 
They are in general very complicated substances, many of them 
occurring naturally in rocks and clays. The chief is aluminium 
silicate, the main constituent of clay, and possibly of the con- 
stitution f~Si(OH)/ \Al(OH)lo or Al 2 Si 2 7 .2H 2 0. 

Sodium metasilicate, Na 2 Si0 3 , made by dissolving silica in 
caustic soda, is soluble in water. The commercial product is 
called water-glass and contains more silica than corresponds to 
the above formula ; it is used in solution as a preservative for 
eggs, since it clogs up the pores of the shell with calcium 
silicate and thus prevents the ingress of air and bacteria. It 
has also been successfully employee! in preserving certain of 
the wooden boats and other objects excavated from the 
Glastonbury Lake Village, etc. 

Glass is a mixture of the silicates of calcium and sodium or 
potassium. Ordinary soda-glass is made by fusing a mixture 
of sand, limestone and soda-ash ; hard-glass is made in a 
similar way except that potash is used instead of soda ; while 
flint glass is made from sand, limestone, litharge and soda or 
potash. Flint-glass is more highly refractive than soda- or 
potash-glass and is therefore used for making lenses ; it is, 


however, soft and easily scratched, therefore microscope 
objectives, etc., should always be cleaned with silk or 
chamois leather and not with the rougher cotton or linen 

Coloured glass is made by adding certain metallic oxides to 
the original mixture or preferably to the glass after fusion. 
Blue glass is made by adding cobalt oxide ; amethyst or purple, 
from manganese dioxide ; red, from purple ofCassius (p. 237) ; 
opalescent , from bone-ash. 

Glass is regarded not as a solid but as a very viscous liquid, 
cooled so far below its freezing-point that it crystallizes only 
very slowly. This phenomenon of " super-cooling " is quite 
common ; all supercooled liquids crystallize after a time and in 
the case of glass this process is called devitrification. Soda- 
glass devitrifies more quickly than the other kinds. Devitri- 
fied glass is brittle and useless. 

Silicon hydrides. Silicon resembles carbon in forming 
many hydrides. These, however, while similar in constitution 
to the corresponding carbon compounds, are very different 
in properties. 

Silicomethane, SiH 4 , is made by the action of hydrochloric 
acid on magnesium silicide (prepared by heating silica with 
magnesium powder) 

Mg 2 Si + 4HC1 = SiH 4 + 2MgCl 2 . 

The gas made in this way is not pure, being mixed with other 
silicon hydrides and hydrogen. Silicomethane is a colourless 
gas which readily takes f^re in the air and as generally pre- 
pared is spontaneously inflammable. The products of 
oxidation are silica and water 

SiH 4 + 20 2 = Si0 2 + 2H 2 0. 

Silicomethane reacts with chlorine to form silicon tetra- 
chloride and hydrochloric acid 

SiH 4 + 4C1 2 = SiCl 4 + 4HC1. 

(Cf. methane, p. 295). 
Silicoethane, Si 2 H 6 , silicopropane t Si 3 H 3 , and silicobutane t 


Si 4 H 10 , are also known. Silicoacetylene, (Si 2 H 2 ) n , was made 
as a yellow crystalline powder in 1900 by BRADLEY. 

Silicon fluoride, SiF 4 , is made by acting upon a mixture 
of silica and calcium fluoride with strong sulphuric acid 

2CaF 2 + 2H 2 S0 4 + Si0 2 = 2CaSO 4 + 2H 2 O + SiF 4 . 
The excess of acid takes up the water formed. 

Silicon fluoride is a colourless fuming gas with a pungent 
smell. It is decomposed by water, forming silicic and hydro- 
fluosilicic acids 

3SiF 4 + 3H 2 = H 2 Si0 3 + 2H 2 SiF 6 . 

The action of hydrofluoric acid upon glass (p. 495) results 
in the formation of silicon fluoride, which is then decomposed 
by the water present. 

Silicon tetrachloride, SiCl 4 , is a colourless fuming liquid 
conveniently made by passing chlorine over an intimate 
mixture of silica and carbon heated strongly in a furnace 
SiO a + 20 + 2C1 2 = Si01 4 + 200. 

It is hydrolysed by water, giving a gelatinous precipitate 
of what may be orthosilicic acid (p. 311) 

Si01 4 + 4H 2 - H 4 Si0 4 + 4HC1. 

Silicochloroform, SiH01 3 , is a colourless mobile liquid 
which fumes in the air and is hydrolysed by water to silicic 
acid at ordinary temperatures, but to silicoformic anhydride, 

>0, at 


2SiHCl, 4- 3H 2 = _ . \0 + 6HC1. 


Silicochloroform is made by passing hydrochloric acid gas 
over heated silicon 

Si + 3HC1 = SiH01 3 + H 2 . 

Silicon carbide, SiC, one particular variety being known 
as carborundum, is made by heating a mixture of sand and 
excess of powdered coke in the electric furnace 
SiO 2 + 30 = SiC + 2CO. 


It is an extremely hard crystalline solid, colourless when pure, 
but black as usually prepared. It is used instead of emery 
for grinding arid polishing, etc. 

Optical Activity. The resemblance of silicon to carbon is 
further illustrated by KEPPING'S discovery (1907) that the 
silicon atom can act as a " centre of asymmetry " and form 
optically active compounds. (See textbooks of organic 


Group in Periodic System : IV ; Symbol : Sn ; Valency : 
2 and 4; Atomic Weight: 118-7; Melting Point: 232; 
Specific Gravity : 7-2. 

History. The use of bronze, which is an alloy of copper 
and tin, by Neolithic man proves that tin must have been 
known in those very remote times. The Hebrew word bedil, 
which occurs in the Old Testament, probably means tin (as in 
Numbers xxxi. 22), although sometimes it is used for any base 
metal, such as lead or tin, which occurs in a silver ore (as in 
Isaiah i. 25) and is separated (badala) from the silver in refine- 
ment. C^SAB (De hello gallico, 5, 12) says, " Nascitur (in 
Britannia) plumbum album in mediis regionibus" PLINY first 
carefully distinguished between this plumbum album (tin) and 
plumbum nigrum (lead). He says that tin was obtained from 
the Cassiterides or Tin Islands (British Isles), the name 
xaaohegos having been given to the metal by the Greeks, 
who possibly derived it from the Arabic qasdlr, tin, although 
the opposite is more likely. 

The apparent similarity between lead and tin was the cause 
of their being known very frequently as the "two leads." 
According to the alchemists, tin was especially under the 
influence of the planet Jupiter and was therefore often called 
by the same name. Thus " GEBER " says " Of Jupiter or Tin. 
We Jignifie to the Sons of Learning, that Tin is a Metallick 
Body, white, not pure [white], livid, and founding little, 
partaking of little Earthinefs ; po/JeJJing in its Root Harfhness, 
Softnefs, and Jwiftness of Liquefaction, without Ignition, and 
not abiding the Cupel, or Cement, but Extensible under the 

TIN 315 

Hammer. . . Its vice is, that it breaks every [metallic] 
Body, but Saturn [lead], and mojt pure Sol [gold] [i.e. when 
alloyed with them it makes them brittle]." 

When sheets of tin are bent, they emit a peculiar noise, 
called the " cry of tin," and hence tin was sometimes called 
by the alchemists the " metallic devil." 

The allotropic modification known as grey tin was first 
noticed by ARISTOTLE. (De mirabilibus auscultationibus, 50.) 

The Latin name stannum was given to the metal about 
1,500 years ago. 

Occurrence. Traces of native tin have been found, but 
the metal generally occurs as the oxide, tinstone. This is 
found in Cornwall, Germany, the East Indies, the Malay 
Peninsula, Australia, China, and South America (Bolivia and 
Peru), as well as in other places. Stream tin is tinstone in 
small granules, found in alluvial deposits. 

Extraction. Tin is prepared by reduction of its oxide 
with carbon in the form of charcoal or anthracite 
Sn0 2 + 2C = Sn + 2CO. 

Tinstone is generally mixed with earthy impurities, pyrites, 
wolfram (FeWO 4 ) and other substances, and therefore has to 
be purified before reduction. The ore is first crushed, and 
levigated in a stream of water, when the earthy matter is 
swept away. The impure stannic oxide is then roasted in a 
current of air, which oxidizes the pyrites, sulphur dioxide and 
arsenious oxide volatilizing away. The calcined ore, which 
contains oxides of iron, etc., as impurities, is again levigatecj, 
when most of the impurities are carried off by the current of 
water and the heavier tin oxide sinks to the bottom. 

The purified tin oxide is then smelted with anthracite in 
small blast-furnaces (e.g. in the Dutch East Indies) or in 
reverberatory furnaces (e.g. in Australia, England, Malay), 

SnO a + 2C = Sn + 2CO. 

The charge for a reverberatory furnace may be as much aa 
4 tons of ore with half a ton to a ton of anthracite. 
The crude tin has to be refined, since it contains more or 


less lead, iron, and arsenic. It is therefore liquated, or heated 
gently until the easily fusible tin [" pojfejjing Jwiftnejs of 
Liquefaction "] melts and can be run off from the liquation- 
dross, consisting of the less fusible metals. The tin is still 
further purified by " poling " or stirring the molten metal 
with green wood (see Copper, p. 220) ; the gases which come 
off from the heated wood carry any impurities to the top, 
where they collect as a scum which can be removed. 

The slag still contains large quantities of tin, and is smelted 
a second time. 

The final purification of tin is effected by electrolysis in 
wooden vats lined with asphalt. The electrolyte is an aqueous 
solution of hydrofluosilicic acid (H 2 SiF 6 ) in which tin has been 

Properties. Tin is a lustrous white metal which melts at 
232 and boils at 2,270. It is stable in moist air at ordinary 
temperatures, but molten tin slowly oxidizes in air, forming 
stannic oxide, Sn0 2 . If heated to whiteness (1,500-1,600) 
tin burns in air with a bright flame 
Sn + O 2 = Sn0 2 . 

It is soluble in hydrochloric acid, dissolving slowly in the 
cold dilute acid but very quickly in the hot concentrated acid, 
forming stannous chloride and hydrogen 
Sn + 2HC1 = SnCl 2 + H 2 . 

Concentrated sulphuric acid dissolves the metal on heating, 
sulphur dioxide being evolved and stannous sulphate left 

Sn + 2H 2 S0 4 = SnS0 4 + 2H 2 O + S0 2 . 

Dilute nitric acid slowly dissolves tin, yielding stannous 
nitrate and ammonium nitrate 

4Sn + 10HN0 3 = 4Sn(NO 8 ) 2 + NH 4 N0 8 + 3H 2 0. 

The concentrated acid, if absolutely pure, has no action on 
tin, but in the presence of a little water a vigorous reaction 
occurs and the tin is converted into metastannic acid, a 
hydrated stannic oxide of variable composition, with evolu- 
tion of volumes of brown fumes. If the residue of meta- 
stannio acid is strongly heated, stannic oxide is left. 

TIN 317 

Caustic alkalis dissolve tin on heating, to form stannates 
and hydrogen, e.g. 

H 2 O + Sn + 2NaOH = Na 2 Sn0 3 + 2H f . 

Sodium stannate. 

Allotropic Forms. Tin exists in three allotropio forms. 
Ordinary tin is called white tin, and is stable between the 
limits of temperature 18 and 170. Below 18, white tin 
gradually passes into grey tin, which is a grey powder. 18 is 
therefore the transition point of these two forms ; at 18 both 
forms are equally stable, above it white tin is the stable form, 
and below it, grey tin. This transformation of white tin into 
a grey powder was noticed by ARISTOTLE and was rediscovered 
by ERDMANN in 1851. It was forcibly brought home to the 
Russian Government in 1867, when a consignment of block 
tin that had been kept in cellars at St. Petersburg (bolshevike 
Leningrad) during the winter was found to have disappeared 
when the cellars were opened in the following spring, the 
floor being covered with a grey powder. The grey form may 
easily be reconverted into ordinary white tin by fusion and 
subsequent solidification. 

The rate of transformation of white tin into grey tin depends 
upon two factors : (i) distance of the temperature of trans- 
formation from the transition point ; (ii) the ordinary effect 
of changes of temperature upon the rate of chemical change. 
Factor (i) shows that the lowering of the temperature below 
18 will accelerate the change, while factor (ii) will act in the 
opposite way. Hence there is a temperature of maximum 
velocity of the change, and this is found experimentally to 
be 50. The change takes place much more quickly in the 
presence of a solution of " pink salt " [(NH 4 ) 2 SnCl 6 , p. 320]. 

If white tin is heated to a temperature of about 170, it 
slowly changes into a third allotropic form, called rhombic tin. 
The exact temperature of the transition point is not known ; 
170 is a mean value. 

The following diagram shows the relationships between the 
three forms 

Grey tin -.-. . |N ordinary tin N rhombic tin. 

Transition Transition point, 

point 18 C. about 160- 170 C. 


USES. Tin finds many important application, since it is 
stable in moist air. Tin-foil is used for wrapping chocolates 
and cigarettes, though a good deal of so-called " tin-foil " nowa- 
days is aluminium foil ; tin-plate is sheet-iron covered with 
tin ; while tin is a constituent of many alloys such as bronze, 
phosphor-bronze, pewter, solder, and Britannia metal. 

For the composition of bronze and phosphor-bronze, see 
p. 227. Pewter consists of 4Sn to IPb and a little anti- 
mony, though some varieties have a smaller proportion of 
tin. Too little tin is undesirable, as lead poisoning may occur 
from drinking from vessels with a high proportion of lead. 
Solders are of many varieties, but usually consist of 34-50 
per cent, of tin and the rest lead. Fine- solder is a mixture 
of 1 part of lead with 2 parts of tin. Britannia metal con- 
tains tin, antimony and copper, while certain antifriction alloys, 
used for bearings, contain the same metals in varying pro- 

Tin-plating is carried out by thoroughly cleaning steel or 
iron plates in dilute hydrochloric acid and then passing them 
into a bath of molten tin preserved from oxidation by a layer 
of oil or grease. The plates are afterwards sent through 
rollers to remove excess of tin and are then left with a bright 
and even coating of the metal. 

Large quantities of tin are used in this way, and since tin 
is an expensive metal, many methods have been suggested 
for recovering it from scrap tin-plate. Electrolytic methods 
are sometimes used, but chlorine is now so cheap that the 
scrap tin-plate is usually treated with this gas, which converts 
the tin into the volatile stannic chloride, SnCl 4 . This can be 
hydrolysed into hydrated stannic oxide, from which the tin 
may be recovered. 


Tin forms two oxides, SnO, stannous oxide, and SnO t , 
stannic oxide, and two corresponding series of salts. In 
stannous salts the metal is bivalent and in stannic salts 

Stannous oxide, SnO, is obtained as a dark green, almost 

TIN 319 

black, powder by heating stannous oxalate out of contact 
with air 

SnC 2 4 = SnO + C0 2 + CO, 

and in other ways. It takes fire if heated in air, forming the 
dioxide, and dissolves in dilute acids to give stannous salts. 

Stannic oxide, SnO a , occurs naturally in the crystalline 
form as tinstone or cassiterite. It is formed as a white powder 
by burning tin in air or by heating the metastannic acid pro- 
duced by the action of concentrated nitric acid on the metal. 

Stannic oxide has practically no basic properties and will 
not dissolve in acids ; it is slightly acidic, and if fused with 
caustic alkalis forms salts, the stannates 

SnO a + 2KOH = K 2 Sn0 3 + H 2 0. 

Many hydrates of stannic oxide are known ; they exhibit 
feebly acidic properties and are called stannic acids. Meta- 
stannic acid, or ft -stannic acid, is made, as stated above, by 
the action on tin of concentrated nitric acid. Its formula is 
not definitely known, but the sodium salt, made by the action 
of cold caustic solution upon the acid, has the formula 
Na 2 Sn 6 O n .4H 2 O, so that the acid may be H 2 Sn 5 O n or 
H 2 Sn 6 O u .4H 2 0. 

a-stannic acid, H 2 Sn0 8 , is obtained as a gelatinous precipi- 
tate by adding dilute hydrochloric acid to a solution of sodium 
a-stannate, Na 2 Sn0 8 , made by dissolving stannic oxide in 
fused caustic soda. 

Sodium a-stannate is used as a mordant in dyeing, under 
the name of " preparing-salt." 

Stannous chloride, SnCl a , is formed when tin is dissolved 
in concentrated hydrochloric acid. By evaporation of the 
solution, white crystals of SnCl 8 .2H 2 are formed 

Sn + 2HC1 = SnCl a + H a . 

The hydrated salt on heating is converted into a basic salt, 
with loss of hydrochloric acid, so that the anhydrous salt 
must be obtained in other ways, e.g. by heating tin in a 
stream of hydrochloric acid gas. 


Stannotis chloride is partially hydrolysed by water to a white 
insoluble basic chloride 

2SnCl 2 + 2H 2 = SnCl 2 .Sn(OH) a 
its solution is always made up therefore in dilute hydro- 
chloric acid, which prevents hydrolysis, and not in pure water. 
A solution of stannous chloride in hydrochloric acid is 
largely used as a powerful reducing agent, owing to the 
readiness with which the tin passes to the stannic state. 
Thus, it reduces ferric salts to ferrous, mercuric salts first to 
mercurous salts and then to mercury, nitric acid to hydroxyl- 
amine, etc. 

SnCl 2 + 2FeCl 3 = SnCl 4 + 2FeCl 2 . 
( (i) SnCl 2 + 2HgCl 2 = SiiCl 4 + Hg 2 Cl 2 . 
t(ii) SnCl 2 + Hg 2 Cl 2 = SnCl 4 + 2Hg. 
4SnCl 2 + 8HC1 + HN0 8 = NH 2 .OH + 4SnCl 4 + 3H 2 0. 
It is used as a mordant, under the name of " tin- salt," and 
also in the preparation of purple of Cassius (p. 237). 

Vapour density determinations show that some of the 
molecules of the vapour are Sn 2 Cl 4 and some SnCl 2 , but 
CASTOBO has shown (1898) that the depression of the freezing- 
point of urethane (an organic compound) caused by solution 
in it of stannous chloride corresponds to the formula SnCI 2 . 
Stannic chloride, SnCl 4 , was discovered in '605 by 
LIBAVIUS, and is still sometimes called spiritus fumans Libavii. 
It is prepared by heating tin in a current of dry chlorine 

Sn + 2C1 2 = SnCl 4 . 

It is a colourless fuming liquid, which readily combines with 
water to form solid hydrates, e.g. SnCl 4 .5H 2 0, which is used 
as a mordant in dyeing and is commercially known as " butter 
of tin." With excess of water stannic chloride is partially 
hydrolysed. It will combine directly with ammonium 
chloride to form (NH 4 ) 2 SnCl 6 , ammonium stannichloride, 
which was formerly used as a mordant for certain red dyes 
and was therefore called pink salt. It is the ammonium salt 
of hydroclilorostannic acid, H 2 SnCl c . 

Stannous sulphide, SnS, is formed when tin is heated with 
sulphur, and when sulphuretted hydrogen is passed through 

LEAD 321 

a solution of stannous chloride in the presence of hydrochloric 

SnCl, + H 2 S ;= SnS + 2HC1. 

It is a brownish-black substance which dissolves in dilute 
acids, forming stannous salts, and in yellow ammonium 
sulphide, forming ammonium thiostannate, (NH 4 ) 2 SnS 8 
SnS + (NH 4 ) 2 S 2 = (NH 4 ) 2 SnS 3 . 

Stannic sulphide, SnS 2 , is made by heating a mixture of 
tin amalgam, sulphur and ammonium chloride, when it is 
left in the form of golden-yellow crystalline plates, called 
mosaic gold. When sulphuretted hydrogen is passed through 
an acidified solution of stannic chloride, stannic sulphide is 
obtained in an impure state as a dirty yellow precipitate, 
soluble in yellow ammonium sulphide, forming ammonium 

Organic Compounds of Tin. The similarity of tin to 
carbon is well shown by the fact that many organic tin com- 
pounds are known, in which the tin is quadrivalent. Thus 
tin ethyl, (C 2 H 6 ) 4 Sn, corresponding to tetra-ethyl methane 

has been prepared, and so has /Sn = O, dimethyl 

stannone, corresponding to acetone, yC O. 

CH 3 / 

OH, \ /CH 2 .CH 2 .CH 3 
The compound ySn.<f has been ob- 

tained in optically active forms by POPE and PEACIIEY (1900). 


Group in Periodic System : IV ; Symbol : Pb ; Valency : 
I, 2, &nd 4 ; Atomicity of Vapour : 1 ; Atomic Weight : 
207-2; Melting Point : 326; Specific Gravity : 11-25-11-4. 

History. Lead has been known from very remote times, 
and was early distinguished as plumbum nigrum from plumbum 
album (tin). The Arabian chemists, while recognizing the 
similarity of the two metals, realized that they were distinct, 
and called lead usrub while they gave the name qasdlr to tin. 



Astrologically lead was connected with the planet Saturn and 
is very often referred to by the name of tliat planet in 
alchemical treatises. " GEBEB " says that " Lead is a 
Metallick Body, livid, earthy, ponderous, mute, partaking of 
a little Whiteness, with much palenejs, refujing the Cineritium 
and Cement, eajily in all its dimenfions with Jmall Compreffion 
extenfible, and readily fujible, without Ignition. . . . Lead is 
in like Manner burnt, and made Minium ; and it is put over 
the Vapours of Vinegar, and made Cerufs. . . . Lead aljo is 
the Tryal of Silver in the Cupel" 

Occurrence. The chief ore of lead is galena, lead sul- 
phide, PbS. It also occurs in the form of carbonate, PbCO,, 
in the mineral cerussite. 

Extraction. Practically all the lead of commerce is made 
from galena, which occurs in England, New South Wales, and 
many other countries. The galena is first roasted in a current 
of air, which converts part of it into lead oxide and part into 

(i) 2PbS + 30 2 = 2PbO + 2S0 2 . 
(ii) PbS + 20 2 = PbS0 4 . 

The roasted ore is then mixed with more galena and the 

temperature raised ; at the same time the current of air is 

cut off. The following reactions then occur 

(i) 2PbO + PbS = 3Pb + S0 2 . 

(ii) PbS0 4 + PbS = 2Pb + 2SO 2 . 

The molten lead collects at the bottom of the furnace and 
is run off. It contains impurities, chiefly antimony, tin, 
copper, and bismuth, which are removed by heating the lead 
to redness in a small reverberatory furnace until the impurities 
are oxidized and collect on the surface of the metal as a scum 
which is removed from time to time. If much copper is present 
this is removed by liquation, i.e. heating the impure lead until 
it melts, when it may be run off from the less fusible copper. 

The lead purified in this way still contains silver, which ia 
often present in sufficiently large quantities to be worth 
extraction. (See desilverization of lead, p. 228.) 

Properties. According to LAMBERT and CULLIS (1915) 

LEAD 323 

absolutely pure lead is a silvery white metal, but as usually 
obtained it has a bluish grey colour. It rapidly tarnishes in 
the air, owing to the formation of a thin film of lead suboxide, 
Pb 2 0. The specific gravity of the metal is 11-3, its melting- 
point 326 and its boiling-point over 1,200. 

Lead is a soft metal, and leaves a mark if rubbed across 
paper. It can be cut with a knife, and is readily malleable. 
It can easily be moulded into pipes or pressed into sheets, and 
as it oxidizes only superficially in air it finds many applications. 
Alloys of lead with antimony and tin are used as type-metal, 
while the " lead " of accumulator plates frequently consists 
of an alloy of 24 parts of lead with 1 of tin. 

It is soluble in nitric acid, forming lead nitrate and nitrogen 
peroxide, etc. 

Pb + 4HN0 8 = Pb(N0 3 ) 2 + 2N0 2 + 2H 2 0, 
or perhaps 3Pb + 8HN0 8 = 3Pb(N0 3 ) 2 + 4H 2 + 2NO. 

Hot concentrated sulphuric acid rapidly dissolves lead, with 
formation of lead sulphate and sulphur dioxide. The dilute 
acid has no action, neither has hydrochloric acid. 

By heating lead tartrate, the metal is obtained in a very 
finely divided condition and is then called pyrophoric lead, 
since it takes fire spontaneously in the air. 

The atomic weight of ordinary lead is 207-2, but lead exists 
in isotopic forms, two, of atomic weights 206 and 208 respect- 
ively, being of considerable interest in that they were the 
first isotopes the existence of which was experimentally proved. 
Lead and its soluble compounds are poisonous, and are 
all the more dangerous in that the effect is cumulative ; i.e. 
the poisonous substances are not eliminated from the system. 
Lead poisoning is common among painters and others who 
work with lead or compounds containing it. 
COMPOUNDS OF LEAD. Lead forms five oxides 
Pb 2 0, lead suboxide. 
PbO, lead monoxide or litharge. 
Pb 3 4 , triplumbic tetroxide, red lead, or minium. 
Pb 2 8 , lead sesqui-oxide. 
PbO 2 , lead peroxide. 


[Several more have been reported, but their individuality 
is doubtful.] 

Lead suboxide and sesqui -oxide are unimportant. The 
former is obtained by gently heating lead oxalate in absence 
of air 

2PbC 2 O 4 = Pb 2 + 3C0 2 + CO, 

and is a black powder. It forms, under suitable conditions, 
the corresponding sub-salts. Lead sesqui-oxide, Pb 2 3 , is a 
reddish yellow powder, discovered by SCHAFFNER (1844), 
who made it by adding sodium hypochlorite to a cold solution 
of litharge in caustic soda. 

Lead monoxide, massicot, or litharge, PbO, is formed 
when lead is strongly heated in air, or by the action of heat 
on lead nitrate 

2Pb(N0 3 ) 2 = 2PbO + 4NO, + 2 . 

It is a yellow powder, commonly known as massicot ; on 
fusion and cooling it is converted into the reddish crystalline 
variety, called litharge, since it is obtained in the purification of 
silver. (Litharge = " silver stone.") Lead monoxide is readily 
reduced to the metal ; it dissolves in acids to give lead salts, 
and also in concentrated caustic alkalis, to form salts, the 
plumbites it is therefore both basic and acidic, although more 
the former than the latter. 

It is used in glass manufacture and as a source of more 
important compounds of lead such as " red lead " and " white 
lead " ; under carefully regulated conditions and in specially 
licensed factories it is used for glazing pottery. (Precautions 
have to be taken since lead compounds are poisonous.) It is 
also used for " drying " paints and varnishes, or making 
them " set " ; this is an oxidation process and therefore 
hastened by the ready " reducibility " of litharge, although 
it is said that the action is also catalytic. 

Litharge is slightly soluble in water, with which it partly 
combines, to form lead hydroxide. The solution has a 
perceptibly alkaline reaction 

PbO + H,0 ^z Pb(OH) a ^zi Pb" + 20H'. 

LEAD 325 

Lead hydroxide, Pb(OH) 2 , is precipitated as a white solid 
on addition of caustic alkali solution to a solution of a soluble 
lead salt. It is both basic and acidic. 

Red lead or minium, Pb 3 4 , has been used for centuries 
as a red paint. It is said that " miniature " really means 
" a picture painted by a ' miniatore '," i.e. one who used red 
lead. Red lead is made by heating litharge in air for some 
hours at a temperature of 400-450. It seems to be a com- 
pound of the basic litharge, PbO, with the acidic lead peroxide, 
Pb0 2 , and is therefore sometimes called lead orthoplumbate, 
the lead salt (Pb 2 Pb0 4 ) of orthoplumbic acid, Pb(OH) 4 or 
H 4 Pb0 4 . It reacts with acids giving ordinary lead salts and 
a deposit of lead peroxide 

Pb 3 4 + 4HN0 3 = 2Pb(N0 3 ) a + 2H 2 + Pb0 8 . 

Lead dioxide or peroxide, Pb0 2 , is a chocolate-brown 
powder obtained by the action of an oxidizing agent, such as 
potassium chlorate, upon litharge, but most conveniently 
prepared from red lead by the reaction given above. It ia 
formed as a brown deposit on the anode when a solution of a 
lead salt is electrolysed between lead electrodes, or by the 
electrolysis of dilute sulphuric acid, using a lead anode. In 
the latter case the S0 4 " ions after discharge at the anode act 
on the water present and liberate oxygen 

2S0 4 + 2H 2 O = 2H 2 S0 4 + 20 ; 

the oxygen then oxidizes the lead anode to lead peroxide 
Pb + 20 = Pb0 2 . 

This reaction is used in accumulators (Fig. 67). A dis- 
charged accumulator consists essentially of two lead plates 
coated, with lead sulphate and immersed in dilute (20 per 
cent.) sulphuric acid. To charge the cell, one of the plates 
is connected to the positive pole of a dynamo, and thus made 
the anode, while the other is made the cathode. On electro- 
lysis, S0 4 " ions pass to the anode, where they are discharged, 
lead peroxide being deposited 

SO 4 + PbS0 4 + 2H a O ^n PbO, + 2H 2 SO A . 


At the cathode, the hydrogen ions are discharged and the 
lead sulphate is reduced to lead 

PbS0 4 + 2H ;zz Pb + H 2 S0 4 . 

Charging is therefore continued until gases begin to appear 
on the plates, showing that the above changes are complete. 
The cell will now give a current (flowing in the opposite 
direction to the charging current) if the two plates are con- 
nected by a wire. The charges that occur during discharge 
are the reverse of those that take place during charging 

Pb0 2 + 2H 2 S0 4 ->PbS0 4 + 2H 2 + S0 4 // at the 

positive pole, 

and Pb + H 2 S0 4 -> PbS0 4 + 2H' at the negative pole. 

After discharge, therefore, both plates are covered with lead 

The principle of the accumulator is thus the conversion 
of electrical energy into chemical energy (charging), followed 
by the re-conversion of this chemical energy into electrical 
energy (discharging). Expressed in one equation the changes 
that take place are 

(i) Charge. 2PbS0 4 + 2H 2 O = Pb0 2 + 2H 2 SO 4 + Pb, 

with absorption of energy, supplied by the 
(ii) Discharge. Pb0 2 + 2H 2 S0 4 + Pb = 2PbS0 4 + 2H 2 0, 

with liberation of energy, given out in the 
form of an electric current. 

Lead peroxide is used as an oxidizing agent ; thus it 
oxidizes hot hydrochloric acid to chlorine 

Pb0 2 + 4HC1 = PbCl 2 + 2H 2 + C1 2 , 

and combines with sulphur dioxide, to form lead sulphate, 
with so much evolution of heat that it becomes incandescent 

A mixture of sulphur and lead peroxide inflames, and a 
mixture of red phosphorus and lead peroxide explodes, when 
rubbed in a mortar. 

Lead peroxide exhibits both basic and acidic properties j 



it dissolves in cold concentrated hydrochloric acid to give a 
dark solution containing lead tetrachloride, PbCl 4 , probably in 
the form of hydrochloroplumbic acid, H[ 2 PbCl fl (-= PbCl 4 .2HCl). 
It also dissolves in hot concentrated caustic alkalis to form 
salts, the plumbates e.g. 

2NaOH + Pb0 2 = Na 2 Pb0 8 + H a O. 
Sodium plumbate. 

Other plumbates are known ; they are the salts of the plumbic 






FIG. 67. 

Pb0 2 .2H 2 0, Pb(OH) 4 , or H 4 Pb0 4 , or^oplumbic acid, 
known in the form of its calcium salt, Ca 2 Pb0 4 ; and 

Pb0 2 .H 2 0, Pb-OH, or H 2 Pb0 3 , mefoplumbic acid, 

which lias been prepared in the free state. 

Lead dichloride, PbCl a , is generally made as a white 
crystalline precipitate by adding a solution of a soluble 
cnloride to a solution of lead nitrate or acetate (the two 
common soluble salts of lead) 

Pb(N0 8 ) a + 2NaCl = PbCl a * + 2NaNO,. 


It may also be made by heating lead in chlorine, since the 
higher chloride, PbCl 4 , which would be expected, dissociates 
on heating. 

Lead chloride, as it is commonly called, is only slightly 
soluble in water, but sufficiently so for a part of the lead in 
qualitative analysis to be left over from Group I and to make 
its appearance as lead sulphide in Group II. 

It is more soluble in hot water than in cold. 

Lead tetrachloride, PbCl 4 , is prepared from the solution 
of hydroch lor opium bic acid described under lead dioxide. 
Ammonium chloride is first added, when a yellow precipitate 
of the ammonium salt, (NH 4 ) 2 PbCl 6 , is obtained. If this 
precipitate is then added to well-cooled concentrated sul- 
phuric acid, hydrochloric acid gas is evolved and a yellow oil, 
lead letrachloride, PbCl 4 , separates 

(NH 4 ) 2 PbCl 6 + H 2 S0 4 - (NH 4 ) 2 S0 4 + 2HC1 + PbCl 4 . 

It fumes in moist air, is readily hydrolysed by water, and 
decomposes on slight warming, into the dichloride and 

Lead carbonate, PbC0 3 , is obtained as a white precipitate 
by adding a solution of sodium bicarbonate to a solution 
of lead nitrate or acetate. Sodium carbonate precipitates 
not the normal lead carbonate, but a basic carbonate, 
2PbCO 3 . Pb(OH) 2 , called white lead. This substance is extens- 
ively used for making paints, and is therefore made on a large 
scale commercially. Many processes are employed, but the 
best is still that described by " GEBEB " in the thirteenth 
century and now called the Dutch Process. Rolls of lead 
sheets are placed in earthenware pots, towards the bottom of 
which are shelves or ledges upon which the lead rests. The part 
of the pots below the ledges is filled with vinegar, and the pots 
then covered with lead plates and placed, to the number of a 
gross or more, in heaps of horse-dung. Fermentation of the 
dung serves two purposes, (i) it keeps the pots rather warm 
(" temperature of putrefaction," one of the " fixed points " 
on the old alchemical scale of temperature), and (ii) it liberates 
carbon dioxide. After a time, which varies from four to 

LEAD 329 

twelve weeks, the lead is converted into white lead of a very 
good (indeed, the best) quality. 
Possible reactions 

(i) The acetic acid vapour, in the presence of air, converts 
the lead into a basic lead acetate 

2Pb + 2CH 3 .COOH + 2 - Pb(OH) 2 .(CH 3 .COO) 2 Pb. 

(ii) The carbon dioxide then reacts with the basic acetate 
to form the basic carbonate and acetic acid 

3Pb(OH) 2 .(CH 3 .COO),Pb + 4CO 2 + 2H 2 

= 2Pb(OH) 2 .2PbC0 3 + 6CH 3 .COOH. 

The liberated acetic acid attacks more lead and so the process 
is continuous as long as any lead is left. Any soluble lead 
acetate remaining in the white lead is removed by washing 
with water. The white lead is then dried by heating under 
reduced pressure, or by mixing it with oil in a " pug-mill," 
when most of the water separates and the white lead remains 
in association with the oil. 

In England, old tannery bark is often used instead of 
horse-dung. White lead is also made by hanging straps of 
lead in a chamber over dilute acetic acid, and blowing in 
carbon dioxide, but attempts to hurry matters in this and 
other ways result in the formation of an inferior quality of the 
product, owing to the particles of white lead being rather 
larger in size and therefore possessing a smaller " covering- 

Lead carbonate as a paint suffers from the defect that it is 
blackened by sulphuretted hydrogen ; the white paint in 
chemical laboratories, etc., is therefore made from zinc 
carbonate, which possesses less covering-power but is un- 
changed in colour by sulphuretted hydrogen, since zinc 
sulphide is white. 

Lead sulphate, PbSO 4 , is a heavy white substance, 
obtained by adding sulphuric acid or a solution of a sulphate 
to a solution of a soluble lead salt. It is soluble in hot con- 
centrated sulphuric acid, but most of it separates out again 
on cooling. It is a common impurity in commercial lead- 
ch amber sulphuric acid. 



Lead nitrate, Pb(N0 3 ) 2 , is a white crystalline solid that 
decrepitates on heating and is decomposed into litharge, 
oxygen and nitrogen peroxide 

2Pb(N0 8 ) 2 - 2PbO + 4N0 2 + 2 . 

Lead acetate or sugar of lead, (CH 3 COO) 2 Pb.3H 2 0, 
is made by dissolving litharge in dilute acetic acid and evapor- 
ating the solution to crystallization. It is a poisonous sub- 
stance with a sweet taste, and is readily soluble in water. 
Filter-paper impregnated with lead acetate smoulders when 
ignited, owing to the formation of the spontaneously-com- 
bustible pyrophoric lead. Basic acetates of lead also exist. 

Lead chr ornate, PbCr0 4 , comes down as a yellow pre- 
cipitate on the addition of potassium chromate to lead acetate 
solution. It is used as a pigment under the name of " chrome 

Organic compounds of lead, such as lead tetraethyl, 
Pb(C 2 H 5 ) 4 , are known. These generally contain quadrivalent 
lead. Lead tetraethyl is the " anti-knock " agent in the now 
familiar " ethyl petrol." 


1. How far is the inclusion of carbon, silicon, tin and lead in a single 
group justified from a consideration of the chemical and physical 
properties of these elements ? 

2. Describe the artificial preparation of the diamond. 

3. How is graphite manufactured ? What are its properties and 
uses ? 

4. Give an account of the manufacture of coal-gas. 

6. Describe the preparation and properties of methane, ethylone 
and acetylene. 

6. How is carbon monoxide prepared in the laboratory ? What 
are its properties ? 

7. Write a short account of gaseous fuels, excluding coal-gas. 

8. Discuss the equilibrium 2NaHCO 8 ^Na 2 CO 3 -f H a O -f CO 2 in 

Solid. Solid. Gas. Gas. 

terms of (a) Le Chatelier's principle, and (6) the phase rule. 

9. A solution of sodium carbonate is alkaline to litmus and a solu- 
tion of sodium bicarbonate is neutral. Explain these phenomena in 
terms of the ionic theory. 

10. Write an account of the silicic acids and silicates. 

11. What do you know of the history of tin T 

LEAD 331 

12. Describe the metallurgy of tin. 

13. Discuss the equilibria between tin allotropes. 

14. Describe the process of tin-plating and state how tin is recovered 
from scrap tin-plate. 

15. How would you prepare (i) anhydrous stannic chloride, (ii) 
anhydrous stannous chloride ? For what purposes is stannous chloride 
used ? 

16. Describe the preparation and properties of the oxides of lead. 

17. Explain the chemlstay of the accumulator. 

18. How is " white lead " made commercially I 


TYPICAL ELEMENTS : Nitrogen, Phosphorus. 

Sub-group A : Vanadium, Niobium, Tantalum. 

Sub-group B (similar to typical elements) : Arsenic, 
Antimony, Bismuth. 


For a group in the Periodic System, which, as we have seen, 
often classifies together elements of widely -diverging character, 
the elements nitrogen, phosphorus, arsenic, antimony and 
bismuth form a remarkably homogeneous family. 

They are all tervalent or quinquevalent, and they show a 
gradual change in physical and chemical properties from the 
characteristic gaseous non-metal nitrogen, through the 
metalloid arsenic, to the true metal bismuth. 


Group in Periodic System : V ; Symbol : N ; Valency : 3 
and 5 ; Atomicity : 2 ; Atomic Weight : 14-01. 

History. Nitrogen was discovered in 1772 by RUTHER- 
FORD, Professor of Botany at Edinburgh University. He 
obtained it from air in the course of experiments on the 
respiration of animals, and called it mephitic air since it would 
not support life. SCHEELE also obtained it from air and 
called it foul air, and it seems to have been independently 
discovered by PRIESTLEY, who gave it the name of phlogisti- 
cated air. LAVOISIER renamed the gas azote, by which name 




it is still known in France, while the name nitrogen was given 
to it on account of the fact that it occurs in nitre. 

Occurrence. In the uncombined state nitrogen forms 
about 79 per cent, by volume and 77 per cent, by weight of 
the air. It is found combined in large quantities as sodium 
nitrate or Chile saltpetre (NaN0 3 ), and is widely distributed 
in smaller quantity in the soil as ammonium salts and nitrates 
of sodium, potassium and calcium. It is an essential con- 
stituent of living matter. 

Preparation. It is difficult to obtain pure nitrogen from 
the air, since although the water- vapour, carbon dioxide, and 
oxygen may readily be removed, the residual nitrogen still 




FIG. 68. 

contains argon and the other rare gases of the atmosphere 
(p. 187), from which it cannot be purified except by long and 
tedious processes, e.g. liquefaction followed by fractional 
distillation. However, for ordinary chemical purposes the 
presence of the inactive gases in nitrogen is no drawback, as 
these substances naturally do not interfere with any of the 
chemical reactions for which the nitrogen may be required. 
It is only when the nitrogen is required for the purpose of 
determining its physical constants, etc., that it is essential 
to obtain it free from all impurities. 

" Atmospheric " nitrogen, then (i.e. nitrogen containing 
about 1 per cent, of the inactive gases), may be prepared by 
removing the oxygen from air which has been previously freed 
from (a) carbon dioxide (by passing through caustic soda 



solution), and (6) water- vapour (by passing through concen- 
trated sulphuric acid). The oxygen may be removed in 
various ways 

(i) By passing the air over heated copper, when copper 
oxide is formed and the nitrogen left. (Fig. 68.) 
(ii) By burning phosphorus in the air. 
It is clear that a great number of similar methods is avail- 
able. Moist nitrogen is readily obtained from air by 
shaking the air up with an alkaline solution of pyrogallol. 
This absorbs both the carbon dioxide and the oxygen, going 

brown in the 
process (owing to 
the formation 
of oxidation pro- 
ducts), and 
leaves the moist 

"Chemical" ni- 
trogen is pre- 
pared by the de- 
composition of 
nitrogenous com- 
pounds, and is 
therefore free 
from the inactive 
gases. (i) The 
usual laboratory 

method is to heat a solution of ammonium nitrite (Fig. 

NH 4 N0 2 = N 2 + 2H 2 0. 

As, however, ammonium nitrite is an unstable compound 
it is never put on the market in the solid state, and even in 
solution is but rarely met with. It is better to use a mixture 
of equimolecular proportions of ammonium chloride and 
sodium nitrite, which in solution react to give ammonium 
nitrite and sodium chloride 

NH 4 C1 + NaNO a ;= NaCl + NH 4 NO t . 

FIG. 69. 


(ii) By passing chlorine into a concentrated solution of 

8NH 3 + 3C1 2 = 6NH 4 C1 + N 2 . 

The chlorine removes hydrogen from some of the ammonia, 
forming hydrochloric acid and nitrogen ; the hydrochloric 
acid then combines with excess of ammonia to form ammo- 
nium chloride. Care must be taken not to pass the chlorine 
for too long a time, or the highly explosive nitrogen trichloride, 
NC1 3 , will be formed. 

(iii) By the action of heat on ammonium dichr ornate 

(NH 4 ) 2 Cr 2 7 = N 2 + Cr 2 3 + 4H 2 0. 

Nitrogen and steam are given off and a green solid, chro- 
mium sesqui-oxide, is left. The reaction is too vigorous for 
convenience unless the dichromate is mixed with about twice 
its volume of dry sand. A little ammonia is formed as a 

Commercially, nitrogen is prepared from the atmosphere by 
the fractional distillation of liquid air. It was formerly a 
waste-product, but is now largely employed in the manufac- 
ture of certain nitrogen compounds (p. 340, etc.). 

Properties. Nitrogen is a colourless, odourless gas. It 
will not burn nor support combustion. It is very slightly 
soluble in water, less so than oxygen. It boils at 194 and 
the liquid nitrogen freezes to a white solid at 214. Under 
ordinary conditions nitrogen behaves as a rather inert element, 
but of recent years much work has been done on the com- 
bination of nitrogen with other elements and the necessary 
conditions discovered. The compounds of nitrogen are very 
numerous, interesting and important. Explosives, dyes, 
drugs; and artificial manures are mostly nitrogenous com- 
pounds, and nitrogen compounds are necessary to the life of 
plants and animals. 

Nitrogen will combine directly with many metais, on 
heating, forming nitrides 

3Mg + N 2 = Mg 3 N 2 , magnesium nitride. 
6Li + N a = 2Li 3 N, lithium nitride. 



Fig. 70 shows an apparatus that may be used for demon- 
strating the absorption of nitrogen by heating magnesium. 
It is interesting to note that when magnesium burns in the 
air it burns in both the oxygen and the nitrogen, the product 
being a mixture of magnesium oxide with a little magnesium 
nitride. This may easily be proved by adding a little water 
and then some Nessler's solution (p. 350). When a metallic 
nitride is acted upon by water, ammonia is formed 

Mg 3 N 2 + 6H 2 = 3Mg(OH) 2 + 2NH 3 , 

and Nessler's solution with ammonia gives a yellowish pre- 
cipitate or colora- 
tion. Magnesium 
nitride is a highly 
magnetic sub- 

Nitrogen and 
oxygen, and nitro- 
gen and hydrogen, 
combine together if 

2N + 2 ;= 2ND 

but the reactions 
are reversible and 
FIG. 70. the equilibrium 

mixtures obtained 

in this way contain only traces of oxides of nitrogen and 
ammonia respectively. However, the fact that a little 
combination does occur is of great importance, as we shall 
gee later. 

Nitrogen will combine directly with many metallic carbides 
if heated, e.g. 

CaC 2 + N 2 = CaCN 2 + C. 

CaCN 2 is calcium cyanamide and is of technical importance 
(p. 248). 


Active Nitrogen is an allotropic form of nitrogen dis- 
covered in 1911 by STEUTT. If an electric discharge is passed 
through nitrogen at low pressure, the gas continues to glow 
after the discharge has stopped. Investigation of the glowing 
gas showed that it was more active than ordinary nitrogen, 
If passed into the vapour of mercury, cadmium, zinc, sulphur, 
etc., the corresponding nitrides are formed, while acetylene and 
ethylene react vigorously with " active " nitrogen forming 
cyanogen, C 2 N 2 , and prussic acid, HON. It is thought that 
in this form of nitrogen the molecules are monatomic, but the 
problem is not yet solved ; it is doubtful, even, whether 
absolutely pure nitrogen can be converted into the " active " 
form, since a trace of oxygen, carbon dioxide, or hydrogen 
sulphide has to be present in the gas before " activity " can 
be obtained. 


With hydrogen, nitrogen forms three compounds 

Ammonia, NH 8 or H 


/ N \ 
H H 

H \ / H 

Hydmzine, N 2 H 4 or >N N<" 

H/ X H 

N s 

Hydrazoic acid, N 3 H or 

N H. 

Of these, the first two are important. 

Ammonia, NH 8 . 

In the form of its compound with hydrochloric acid, 
ammonium chloride, NH 4 C1, ammonia has been known for 
over 2,000 years. This salt occurs naturally in Armenia and in 
other parts of Asia, and appears to have been called origin- 
ally sal armeniac. The early chemists, however, confused 


it with the natural sodium sesquicarbonate or natron found in 
the Libyan desert near the temple of JUPITER AMMON, and 
the name was changed in the course of time into sal ammoniac 
through this misconception. 

PLINY (Hist. Nat., xxxi, chap. 39) ascribes its discovery to 
KING PTOLEMY " Bang Ptolemy discovered it near Pelusium 
while making a camp. Following his example, men found it 
by digging in the sand in the waste places from Egypt to 
Arabia, and also in the dry localities throughout Africa as far 
as the Oracle of Amnion. It grows by night, according to the 
phases of the moon. The country near Cyrene is famous for 
its sal ammoniac, so called because it is found beneath the 
sand (Greek ammos). ... It grows in long pieces." 

" GEBER," in the thirteenth century, describes its purifica- 
tion in the following words " Grind it first . . . then let it 
be Jublimed in an high Body and Head, until it all ajcend pure. 
Afterward dijjolve it upon a Porphiry in the open Air, if you 
would of it make Water ; or keep the Sublimate Jufficiently 

PRIESTLEY was the first to prepare ammonia gas (1774). 
He made it by heating a mixture of lime and sal ammoniac, 
and collected the gas over mercury. He called it alkaline air. 

Ammonia occurs in the soil in small quantities, and in 
larger quantities, as ammonium sulphate (NH 4 ) 2 S04, in the 
soffioni or fumaroles of Tuscany. It is obtained by the de- 
structive distillation of many organic materials, such as 
horns, hoofs, bones, and coal. The old name for an aqueous 
solution of ammonia, spirit of hartshorn, indicates the method 
of preparation. 

Ammonia is prepared in the laboratory by Priestley's 

Ca(OH) 3 + 2NH 4 C1 = CaCl 2 + 2NH 3 + 2H 2 0. 

It cannot be dried by means of sulphuric acid, phosphorus 
pentoxide, or calcium chloride, as it combines with these 
substances ; it is therefore dried by passing through a glass 
" tower " containing lumps of quicklime, and is then collected 
by upward displacement or over mercury. 



Technical Preparation. Ammonia and its compounds are 

very important commercial products. The chief sources are 

(i) The aqueous liquid obtained in the manufacture of 

(ii) A similar liquid obtained in the manufacture of coke. 

(iii) The distillation of shale. 

(iv) The manufacture of water-gas (p. 301). 
(v) Synthetic and electrical methods. 

(i), (ii) and (iv). Coal contains about 1-5 per cent, of nitro- 
gen, and much of this is obtained in the form of ammonia when 
coal is heated in the absence 
of air. On passing the gases 
through water the ammonia 
dissolves, giving the so-called 
ammoniacal liquor. About 5 Ib. 
of ammonia are obtained from 
every ton of coal in the manu- 
facture of coal-gas, and the 
coke left in the retorts still 
contains much nitrogen which 
is liberated as ammonia when 
steam is passed over the red- 
hot coke in the manufacture of 

The ammoniacal liquors are 
treated with lime and then 
distilled ; the ammonia which 

comes oif is collected in aqueous sulphuric acid and the solu- 
tion of ammonium sulphate so obtained evaporated to crys- 
tallization. The product is an impure ammonium sulphate 
(95 per cent.). Ammonia may be liberated from it, if de- 
sired, by heating with lime ; if the gas is then passed into 
water ammonia solution is obtained. Pure ammonia is 
sometimes made from the crude ammonium sulphate by 
heating the latter to about 200 for some time, to decompose 
organic matter present as impurity, and then raising the 
temperature to 380, when pure ammonia is given off and a 
compound (NH 4 ) 2 S 2 7 left. 

FIG. 71. Preparation of moist 


(iii) Shale. Shale consists of a hard slate-like clay, often 
containing a considerable amount of carbonaceous matter 
(" bituminous shale "), When bituminous shale is destruc- 
tively distilled, an ammoniacal distillate is obtained, from 
which the ammonia can be extracted in the form of 
ammonium sulphate by addition of sulphuric acid followed by 
evaporation of the solution. About 70,000 tons of ammon- 
ium sulphate are prepared annually in this way, chiefly in 

(v) Synthetic methods. When nitrogen and hydrogen are 
sparked together, in the proportions by volume in which they 
occur in ammonia, about 2 per cent, of the mixture is con- 
verted into ammonia, heat being evolved in the process. 
Ammonia is thus an exothermic substance, and we might 
expect that the reaction 

2N + 3H 2 ->2NH 3 

would go much more nearly to completion than is represented 
by the 2 per cent, of ammonia actually obtained. The small 
yield of ammonia can be explained by supposing that the rate 
of formation of ammonia is very small at ordinary tempera- 
tures ; at the temperature of the electric spark the rate would 
certainly be tremendously increased, but so would the rate of 
decomposition of the ammom'a formed. Hence to synthesize 
ammonia successfully from nitrogen and hydrogen it seems 
that a suitable, catalyst is the first thing necessary. 

There are, however, other points to be considered, (i) As 
ammonia is an exothermic compound, increase of temperature 
would lead to a diminution of the proportion of ammonia in 
the equilibrium mixture of this gas with nitrogen and hydro- 
gen, (ii) The equation N 2 + 3H 2 - ^ 2NH 3 shows us that 
a decrease in volume takes place when ammonia is formed 
from nitrogen and hydrogen. We should therefore expect, 
by Le Chatelier's principle, that increase of pressure would 
aid the formation of ammonia. 

To prepare ammonia synthetically, then, on a scale which 
will assure commercial success, the following conditions have 
to be observed 



(a) The reaction velocity of the combination of nitrogen 
and hydrogen must be increased by raising the temperature 
or by addition of a suitable catalyst, or both. 

(6) The temperature must not be so high that the propor- 
tion of ammonia in the equilibrium mixture is lowered too far 
for the process to be commercially successful. 

(c) High pressures should be employed, to increase the 
proportion of ammonia formed. 

(d) It is clear that the most effective way of preventing 
the reverse reaction (decomposition of ammonia) is to remove 
the ammonia from the sphere of action as soon as it is 

In the following table are shown the percentages of ammonia 
in the equilibrium mixture at various temperatures and 


Pressure In 





G00. ; 650". 







0-08 0-05 





















The catalyst generally used (!!ABEB process) is a mixture 
of finely divided iron and molybdenum. The nitrogen and 
hydrogen must be carefully purified before use ; they are then 
compressed to about 200 atmospheres and passed over the 
catalyst at a temperature of 500. The ammonia so formed 
is cooled and dissolved in water or else liquefied. This process 
is a great commercial success, and as nitric acid may be 
obtained by oxidation of ammonia, this and other methods 
rendered Germany independent of Chile saltpetre during the 
European war of 1914-1918. 



Cyanamide process. When calcium cyanamide, CaCN, 
(p. 248), is treated with steam, ammonia is formed 

CaCN 2 + 3H 2 - CaC0 3 + 2NH 8 . 

As calcium cyanamide is now prepared on a large scale by 
passing nitrogen over calcium carbide heated in the electric 
furnace, the preparation of ammonia in this way is used for 
commercial purposes. 

Properties. Ammonia is a colourless gas with a pungent 

FIG. 72. Ammonia Ice-making Machine. 

A. The compressed ammonia Is cooled in A (through which cold water circulates) and 
liquefies. B. In coil B the liquid ammonia is under reduced pressure and evaporates 
rapidly, the gas passing back to the pump whence it passes to A again, and so on. C. 
Vessel containing calcium chloride solution which does not freeze. Z), D. Vessels con- 
taining water which is frozen. E, E. Stirrers. 

Bmell. It is lighter than air and therefore may be collected 
by upward displacement, especially as it is extremely soluble 
in water and therefore cannot be collected at the pneumatic 
trough. Ammonia can easily be liquefied, as was first shown 
by FARADAY. Liquid ammonia is a colourless liquid boiling 
at 33*5 and solidifying to a colourless crystalline solid at 
78. It is used in commerce for refrigerating or ice-making 



since by rapidly evaporating it low temperatures are pro- 
duced ; the gaseous ammonia may then be liquefied again by 
passing it through a condenser surrounded by cold water and 
compressing it. (Fig. 72.) 

Ammonia is not a base, since it combines directly with 
acids to form salts without elimination of water ; it is, in fact, 
the anhydride of a base, the true base being ammonium 
hydroxide, NH 4 OH, which is formed when ammonia is dissolved 
in water. The following equations illustrate the point 

NH 3 + HC1 = NH 4 C1 

basic anhydride + acid = salt but no 


NH 4 OH + HC1 = NH 4 C1 + H 2 O 
base + acid = salt + water. 


action on 
turns red 

Dry ammonia has no 
litmus, moist ammonia 
litmus blue. 

Ammonia attempts to burn in the 
air but cannot quite manage it ; it 
will, however, readily burn in oxygen- 
forming nitrogen, steam, and ammo- 
nium nitrate and nitrite. Ammonia 
is readily oxidized by oxygen in the 
presence of a platinum spiral. The 
heat evolved in the process is sufficient 
to keep the platinum red hot and 
often to ignite the mixture of ammo- 
nia and oxygen. The ammonia 



FIG. 73. Ammonia 
burning in Oxygen. 

flame is of a characteristic brownish yellow colour. 

When ammonia is dissolved in water an evolution of heat 
occurs. H air is blown through the solution, the ammonia is 
rapidly driven off and heat therefore absorbed. This reaction 
was at one time made use of in ice-making (CARRE'S process, 
now superseded by that previously described). The aqueous 
solution of ammonia contains ammonia and also ammonium 
hydroxide, NH 4 OH. This is a weak base and partially 
ionizes into NH 4 * and OH', ammonium and hydroxyl ions. 



If, therefore, ammonia is in contact with its saturated solu- 
tion, the following equilibria occur 

NH 3 


NH 3 

NH 4 OH ^ NH 4 + OH'. 

Pure ammonium hydroxide, NH 4 OH, was not isolated until 
1909, when it was prepared by RUPERT, who obtained it as 
a colourless crystalline solid melting at 79 and rapidly 
decomposing on rise of temperature. 

Ammonium. The group of atoms NH 4 behaves in many 
respects like an atom of an alkali metal such as sodium or 
potassium. The following table shows certain of the resem- 



Chloride . 

NH 4 C1, white crystals 

NaCl, white crystals 

Nitrate . 

NH 4 NO 3 

NaNO 3 


(NH 4 ) a S0 4 

Na 2 SO 4 

Hydroxide . 

NH 4 OH 



(NH 4 ) 2 CO, 

Na 2 CO, 

The salts of the NH 4 group are very familiar in appearance 
and properties to those of sodium, and for this reason the name 
ammonium was given to this group of atoms or radical, to 
indicate its apparent metallic nature. The ammonium group 
is univalent, as the nitrogen atom in it is quinquevalent 
and only four of these valencies are used to attach the 
hydrogen atoms 


H \ I 

M\ -. It is therefore capable of 



combining directly with a univalent acid residue or of 
replacing an atom of sodium, etc., in a compound. Up to the 
present, there is no satisfactory evidence that the ammonium 
radical has been isolated, though claims to this effect have 


been made. It was at one time supposed that if it could be 
isolated it would present a metallic form. If a concentrated 
solution of ammonium chloride is poured over some sodium 
amalgam, the latter swells up and forms a peculiar mass called 
ammonium amalgam, which was considered to be an amalgam 
of mercury with ammonium, NH 4 . On standing, the sub- 
stance loses ammonia and hydrogen, and mercury is left. The 
question whether ammonium amalgam is really what it pro- 
fesses to be or not is still unsolved ; there is evidence both ways. 
In any case, the ammonium ion, NH 4 , certainly exists in solu- 
tion. It would be very interesting if " ammonium " itself 
were isolated and proved to have metallic properties ! 

AMMONIUM SALTS. When ammonia combines directly with 
acids the nitrogen atom of the ammonia molecule increases 
in valency from 3 to 5, and ammonium SALTS are formed 
H H 

I H 

N + HC1 = 

H H H 

Chemists were at one time inclined to regard these am- 
monium salts as loose compounds of ammonia with the acid, 
the ammonia and acid molecules being held together by a 
sort of residual or weak valency just as, in cases of hydrated 
salts, the water of crystallization is loosely held. VICTOR 
MEYER, however, in 1874 showed that in ammonium salts the 
nitrogen atom was quinquevalent, in the following way. He 
took trimethylaraine, N(CH 8 ) 3 , and ethyl iodide, C 2 H 5 I, and 
allowed them to react together, thus obtaining ethyltri- 
methylammonium iodide. If this is a loose " molecular com- 
pound " it would have a constitution represented by the 
formula N(CH 3 ) 3 .C 2 H 6 I, whereas if the nitrogen in it is 

CH 3 

pentavalent the formula would be \N I. He next 

C 2 H S 

took ethyldimethylamine, N(CH,) 2 C 2 H 5> and allowed it to 


react with methyl iodide, CH 3 I, obtaining a compound 
identical with the ethyltrimethylammonium iodide previously 
made. Now if the nitrogen is quinquevalent it is clear that 
the reaction of N(CH 3 ) 3 with C 2 H 5 I would yield the same com- 
pound as that produced by the reaction of N(CH 3 ) 2 C 2 H 5 and 


CH 3N | 

CH 3 I, namely ;>N I, whereas if only the loose " mole- 
CH/ | 
C 2 H 5 

cular " compound were formed the first reaction would give 
N(CH 3 ) 8 .C 2 H 5 I and the second N(CH 3 ) 2 C 2 H 5 .CH 3 I, which 
would presumably be different from one another in properties. 
Hence the quinquevalency of nitrogen in the ammonium com- 
pounds may be assumed. Stereochemical work on nitrogen 
compounds of the type NRR'R"R'"I (where R, R', R" and 
R'" are four different carbon radicals) has supported this 

Ammonium hydroxide, NII 4 OH, has been described 
above. It is largely used in qualitative analysis, for the 
precipitation of metals as hydroxides. It is a weak base, that 
is, it is only slightly ionized in aqueous solution, and as this 
ionization may be largely reduced by the addition of am- 
monium chloride it is possible to cause the precipitation of the 
hydroxides of iron, chromium, and aluminium from a solu- 
tion and at the same time to prevent the precipitation of the 
hydroxides of zinc, manganese, cobalt and nickel even though 
these metals may be present in the same solution. This is the 
principle of the separation of Group III metals from those of 
Group IV (analytical table groups, not periodic system groups) 
(see p. 135). 

Ammonium chloride, NH 4 C1, is a white crystalline solid. 
It is made by boiling ammonium sulphate with sodium 
chloride in aqueous solution 

(NH 4 ) 2 S0 4 + 2NaCl ^r Na 2 S0 4 + 2NH 4 CL 
On concentrating the solution the sodium sulphate crystal- 


lizes out first and may be removed ; the solution of ammonium 
chloride may then be evaporated. It is purified by sublim- 
ation as described by GEBER. It is used in Leclanche cells 
and also in dry cells. Its vapour is dissociated (p. 82). 
Ammonium chloride is used in soldering (as a " flux "), 
because the hydrochloric acid which is set free on heating 
cleans the surface of the metal and thus enables the solder to 
" bite." It is also used in the manufacture of dyes and in 
cah* co-printing . 

Ammonium sulphate, (NH 4 ) 2 S0 4 , is obtained from 
ammoniacal liquor (p. 339). It is a white crystalline solid 
very largely used as an artificial manure. It is also used in 
the fermentation industries. It is decomposed on heating, 
one of the products being ammonia, while the other is am- 
monium hydrogen sulphate, NH 4 HS0 4 . At a higher tempera- 
ture further decomposition occurs, yielding steam, sulphur 
dioxide and nitrogen. 

Ammonium nitrate, NH 4 N0 3 , is a white deliquescent 
crystalline solid, very soluble in water, the process of solution 
being accompanied by absorption of heat ; it is therefore used 
in refrigeration on a small scale, e.g. preparation of ice-cream 
(though here it has now been largely replaced by solid carbon 
dioxide ; see p. 305). It is prepared by passing ammonia 
into fairly concentrated (60 per cent.) nitric acid. If heated, 
it splits up into nitrous oxide, N 2 0, and water 
NH 4 N0 3 - N 2 + 2H 2 0. 

This decomposition occasionally becomes very violent and 
an explosion may result. Ammonium nitrate is an endo- 
thermic substance, and as it is also a powerful oxidizing agent 
it is largely used as a constituent of explosives. A mixture of 
ammonium nitrate and aluminium powder is called " am- 
monal," and is one of the explosives used in Mills' bombs. 
" Amatol " is a mixture of 4 parts of ammonium nitrate with 
1 part of trinitrotoluene (" T.N.T."). Ammonium nitrate 
explosives are usually fairly safe to handle. 

Ammonium carbonate, (NH 4 ) 2 C0 8 , is made commercially 
by heating ammonium sulphate with powdered limestone j 


the sublimate obtained is crude ammonium carbonate. It 
contains also ammonium bicarbonate,, NH 4 HC0 3 % , and am- 

monium carbamate, O = C<^ , the ammonium salt of a 

peculiar acid called carbamic acid, = CS . If com- 


mercial ammonium carbonate is treated with a concentrated 
solution of ammonia it is converted into the normal carbonate, 
a white crystalline solid. 

Ammonium sulphide. Normal ammonium sulphide, 
(NH 4 ) 2 S, is formed when a mixture of two volumes of am- 
monia and one of sulphuretted hydrogen, H 2 S, is cooled to 
18. It is a white crystalline solid, readily decomposing 
even in solution, into ammonia and ammonium hydrogen 
sulphide, NH 4 HS 

(NH 4 ) 2 S - NH 3 + NH 4 HS. 

The latter compound may also be formed by passing 
sulphuretted hydrogen into a strong solution of ammonia 
until no more will dissolve. 

The solution of the ammonium hydrogen sulphide so 
obtained is colourless when freshly prepared, but slowly goes 
yellow on exposure to air. The same yellow solution may be 
made by adding sulphur to the warm colourless solution of 
ammonium hydrosulphide. It contains ammonium poly- 
sulphides, (NH 4 ) 2 S a ., of which the chief is said to be (NH 4 ) 2 S 4 . 
The yellow colour assumed by the colourless solution of 
ammonium hydrogen sulphide on exposure to air is therefore 
to be explained by (i) oxidation, with liberation of sulphur, 
followed by (ii) solution of the sulphur in excess of ammonium 
hydrogen sulphide to form ammonium polysulphides. 

" Yellow ammonium sulphide " (as it is called) is used in 
qualitative analysis and also as a reducing agent in organic 

Composition of Ammonia. The vapour density of 
ammonia is 8'5, hence the molecular weight is 17. As the 


atomic weight of nitrogen is 14, the formula must be NH,. 
Additional evidence can be obtained in the following ways 
(i) A measured volume of ammonia is placed in a eudio- 
meter tube over mercury and sparked until no further change 
in volume occurs. The gas is now practically completely 
(98 per cent.) decomposed into nitrogen and hydrogen. The 
volume is noted, and a measured excess of oxygen introduced 
and a spark passed. Explosion occurs, all the hydrogen 
present being converted into steam, which condenses to liquid 
water, the volume of which is negligible compared with the 
volumes of the gases in the experiment. After cooling, the 
residual volume is noted. From the results, a formula for 
ammonia may be calculated. Example 

Volume of ammonia taken = 15 c.c. 

Volume of mixture of nitrogen and hydrogen formed = 30 c.c. 

Volume after addition of oxygen = 58 c.c. 

Volume after explosion 24-25 c.c. 

.*. Contraction on explosion == 33-75 c.c. 

Of this, f will be hydrogen, since 2 volumes of hydrogen combine 
with 1 volume of oxygen to form water. 

i.e. 22-5 c.c. 

But volume of nitrogen -f- hydrogen = 30 c.c. 
.*. Volume of nitrogen = 30 22-5 = 7-5 c.c. 
.*. 15 c.c. of ammonia yield 7-5 c.c. nitrogen and 22-5 c.c. hydrogen. 

2 volumes ,, ,, 1 volume ,, ,,3 volumes ,, 

/. by Avogadro, 2 molecules of ammonia yield 1 molecule nitrogen 

and 3 molecules hydrogen. 
.*. 2 molecules of ammonia consist of 2 atoms of nitrogen and 6 of 


.'. 1 molecule of ammonia contains 1 atom of nitrogen and 3 of hydro- 
gen, and the formula is NH 3 . 

(ii) HOFMAKN'S METHOD. A long tube (Fig. 74), fitted at 
each end with a tap and with a funnel attached at one end, is 
filled with chlorine. In the funnel is placed concentrated 
ammonia solution. On this entering the tube drop by drop, 
vigorous reaction occurs and the tube becomes filled with white 
fumes of ammonium chloride. After the reaction is complete 
the top tap is turned off, and the tube is placed in a pneumatic 
trough containing water. The bottom tap is then opened, 
when water rushes in, and after levelling, the residual gas is 
found to occupy one-third of the volume of the original 
chlorine. This residual gas proves to be nitrogen. The 


hydrogen which was combined with this nitrogen in ammonia 
has been removed by the chlorine ; now chlorine combines 
with its own volume of hydrogen, therefore the volume of 
hydrogen which was combined with the residual nitrogen is 
three times the volume of the latter. In other words, the 
ammonia consists of nitrogen and hydrogen combined 
together in the proportion by volume of 1 to 3. Hence, by 
Avogadro's Hypothesis, the molecule of ammonia 
must be (NH 3 )# ; in other words, the empirical 
formula of ammonia is NH 3 . The true formula can- 
not be determined by Hofmann's method, which 
is therefore not so good as that given under (i). 

x may of course be found by a vapour density 

Tests for Ammonia and Ammonium Salts. All 
ammonium salts when heated with caustic soda 
solution evolve ammonia, which may be detected 
by the smell and by its action on moist red litmus. 
NESSLER'S solution gives a yellow precipitate or 
coloration with an alkaline solution of an ammo- 
nium salt. Nessler's solution is a solution of 
potassium mercuri-iodide, K 2 HgI 4 , made by add- 
ing potassium iodide solution to a solution of 
mercuric chloride till the first-formed precipitate 
of mercuric iodide is just redissolved 

2KI + HgCl 2 = 2KC1 + HgI 2 . 
HgI 2 + 2KI -K 2 HgI 4 . 

The yellow precipitate given with ammonia 
is NHg 2 I. This test for ammonia is used in 
water-analysis. Under suitable conditions it will detect 1 
part of ammonia in 2,000,000 of water. 

Hydrazine, N 2 H 4 , is prepared by the action of sodium 
hypochlorite upon ammonia solution in the presence of glue, 
which acts as a negative catalyst on the decomposition of the 
hydrazine so formed 

NaOCl + NH 3 = NaOH + NH 2 CL 
NH 2 C1 + NH 3 = N 2 H 4 + HC1. 


It is a colourless crystalline solid (M.P. about 1 C.) with a 
weakly basic reaction in aqueous solution. Its derivatives, 
especially phenylhydrazine, C 6 H 6 .NH.NH a , are more impor- 
tant than the base itself. 

Hydrazoic acid, N 3 H, may be obtained as follows. 
Sodamide, Na.NH 2 , is first made by heating sodium in dry 
ammonia. This is then heated in nitrous oxide, when sodium 
hydrazoate is formed 

NaNH 2 + N 2 - H 2 + NaN 3 . 

The sodium hydrazoate is now decomposed with dilute 
sulphuric acid and the liquid distilled, when an aqueous solu- 
tion of hydrazoic acid collects in the receiver. 

Pure hydrazoic acid is a colourless, volatile and very 
explosive liquid, boiling at 37. The lead salt, Pb(N 3 ) 2 , has 
been used as a detonator, but is not so useful for this purpose 
as the more generally employed mercury fulminate. 

Hydroxylamine, NH 2 .OH. This is a colourless deli- 
quescent crystalline solid, which may be obtained by reducing 
nitric oxide with tin and hydrochloric acid, by the electrolytic 
reduction of nitric acid (BOEHRINGER, TAFEL and others), and 
by the action of sulphur dioxide upon a solution of sodium 
carbonate and sodium nitrite. It is a powerful reducing agent 
and is extensively used in organic chemistry, since it gives 
beautifully crystalline compounds called oximes with alde- 
hydes, ketones, and other substances containing the carbonyl 
group, ]>CO. Example 

CH 3 CH 8 

>CO + H 2 N.OH = >C = N.OH + H 2 0, 

CH 3 CH 3 

Acetone., Acetone oxiine 

or acetoxime. 

Hydroxylamine is a basic substance and readily forms salts 
by addition. It is usually kept in the laboratory in the fora) 
of its hydrochloride, NH 2 OH.HC1. 
Oxides of Nitrogen. Nitrogen forms five oxides 

N 2 0, nitrous oxide. 

NO, nitric oxide. 


N 2 3 , nitrogen trioxide or nitrous anhydride. 
N 2 4 or NO 2 , nitrogen peroxide or tetroxide. 
N 2 O 5 , nitrogen pentoxide or nitric anhydride, 
and four oxyacids 

H 2 N 2 2 , hyponitrous acid. 
H 2 N 2 3 , nitrohydroxylaminic acid, 
HN0 2 , nitrous acid. 
HN0 3 , nitric acid. 

Since the starting-point for the preparation of all the rest 
of these compounds is nitric acid, it is convenient to study 
this substance first. 
Nitric Acid, HNO 3 . 

Nitric acid was probably known to the Arabian chemists 

FIG. 75. Preparation of Nitric Acid. 

of the eighth century ; it was certainly prepared by " GEBER " 
in the thirteenth century, who gives detailed instructions for 
the operation 

" First B of vitriol of Cyprus, lib. 1. of Saltpeter, lib. 1 J, and 
of Jamenous Allom one fourth part ; extract the Water with 
Kednefs of the Alembeck (for it is very Solutive) and vje it in 
the before alleadged CJuipters. This is aljo made much more 
acute, if in it you Jhall dijjolve a fourth part of Salammoniac ; 
becauje that dijjolves Gold, Sulphur, and Silver." (The 
" more acute " liquid is of course aqua regia.) l 

The first chemist to make nitric acid by the action of 
sulphuric acid upon potassium nitrate was GLAUBER (1650) 
1 i.e., a mixture of nitric and hydrochloric acids. 



LAVOISIER in 1776 showed that it contained oxygen, and 
CAVENDISH (1784) proved that it contained hydrogen and 
nitrogen as well. 

'Preparation. Nitric acid is prepared in the laboratory, 
and also commercially, by heating potassium or sodium 
nitrate with concentrated sulphuric acid 

(i) KN0 3 + H 2 S0 4 = KHS0 4 + HN0 3 , 
and, on further heating, 

(ii) KHS0 4 + KN0 3 =* K 2 S0 4 + HN0 3 . 

In practice, the action is never taken beyond the first stage, 

FIG, 76. Manufacture of Nitric Acid. 
JL. Entrance for furnace gases. B. Exit for furnace gases. C. Iron retort. D. Receiver*. 

as the temperature required for the second is so high that 
much of the nitric acid is decomposed. 

The distillate consists of nitric acid mixed with a little 
water, which may be removed by addition of concentrated 
sulphuric acid followed by redistillation. The acid so 
obtained is of a yellow colour ; this is due to the presence in 
it of the yellow gas nitrogen peroxide. If a current of dry 
air or dry carbon dioxide is blown through the acid the 
nitrogen peroxide is swept away and the resulting acid is 



Manufacture. Much nitric acid is ma$e, as already 
described, by heating sodium nitrate (Chile saltpetre) with 
concentrated sulphuric acid. However, the supplies of 
Chilean nitrate will eventually be exhausted, and in view of 
this fact, attempts were made to convert atmospheric nitrogen 
into nitric acid. For many years no appreciable advance was 
made toward the solution of this problem, but it has now been 

successfully solved. It is, 
indeed, reported that 
Germany had decided on 
war in 1913, but was forced 
to wait until 1914 becaupa 
her chemists were not 
quite satisfied with tha 
methods for converting 
nitrogen from the air into 
nitric acid. Once these 
methods were perfected, 
Germany became in- 
dependent of Chilean ni- 
trate, and therefore had 
no fear of a shortage of 
explosives or artificial 

The methods of manu- 
facture of nitric acid from 
the air are of two sorts 
(i) Direct combination 
of nitrogen and oxygen 
to form oxides of nitrogen 

FIG. 77. Birkeland Eyde Furnace. 

A. Core. B. Winding. 0. Gas entrance. 
D. Gas exit. 

at the temperature of the electric arc ; nitric acid is then 
made by dissolving these oxides in water. 

(ii) Oxidation of synthetic ammonia (q.v.) to oxides of 
nitrogen and preparation of the nitric acid from these 
oxides as in (i). 

(i) The first successful method of converting the oxygen 
and nitrogen of the air into oxides of nitrogen on a commercial 
ecale was invented by the Scandinavian scientists BIRKB- 


LAND and EYDB in 1903. They pass air through an electric 
arc which is drawn out into a thin disc of flame by means of a 
powerful electromagnet. The disc of arc is some 6-8 feet 
in diameter, and is enclosed in an iron furnace lined with 
resistant material. The issuing gas contains about 2 per 
cent, of nitric oxide, NO. The hot gases are cooled by 
passing through pipes in a boiler, and the steam that is thus 
produced is used to concentrate the solution of nitrates 
obtained in the process. During the cooling the nitric oxide 
present combines with oxygen of the excess of air to form 
nitrogen peroxide (about 600 C.). The gas is then passed 
up an iron or granite tower down which weter is sprayed, 
where part of the nitrogen peroxide is converted into nitric 
acid and part into nitric oxide 

(i) 2N0 2 + H 2 = HN0 3 + HNO a . 
(ii) 2HNO 2 = H 2 + NO + N0 2 . 

The nitric oxide is immediately converted into nitrogen per- 
oxide again by the excess of air present, and thus by having a 
sufficiently large number of condensing towers practically all 
the nitrogen peroxide originally present is converted into nitric 
acid. The solution of nitric acid that collects at the base of 
the towers is pumped up and sprayed in again until it reaches a 
concentration of about 35 per cent. It is then neutralized by 
lime and converted into calcium nitrate, which can be obtained, 
by evaporation of the solution, as a white solid. As normal 
calcium nitrate is deliquescent, it is troublesome to transport 
and handle, so that it is usually converted into a non- 
deliquescent basic calcium nitrate (Norwegian saltpetre or air 
saltpetre) lay addition of a suitable weight of lime. 

" Norwegian saltpetre " is largely used as an artificial 
manure ; for many purposes it is better than sodium nitrate, 
as the lime it contains very often improves the tilth of the soil. 

Sodium nitrite, NaNO 2 , is obtained as a by-product in the 
above process, by absorbing the residual oxides of nitrogen 
in a further tower down which a cool dilute solution of sodium 
hydroxide or carbonate trickles. The commercial success of 
Birkeland and Eyde's process depends upon cheap electrical 


power such as is supplied by the numerous 'waterfalls of 

In the Badische Process, worked at Ludwigshafen, the arc 
is drawn out into a thin flame some 30 feet in length. Special 
forms of apparatus are used to remove the nitric oxide as 
soon as it is formed and to cool it rapidly. 

(ii) That nitric acid may be made by the oxidation of 
ammonia has been known since 1788, arid the French appear 
to have made it in this way for munition purposes during the 
Peninsular Wars, but the process had no commercial value 
until quite recently, after the problem of making synthetic 
ammonia had been successfully solved (p. 3iO). The reaction 
which takes place may be represented by the equation 

4NH 3 + 50 2 = 4NO + 6H 2 0, 

the nitric oxide thus formed being converted into nitric acid 
as described under (i). The final result may be expressed 

NH 3 + 20 2 - HN0 3 + H 2 0. 

for successful oxidation of ammonia by air, the gases must 
be carefully purified and passed over a suitable catalyst 
(generally platinum gauze) at a temperature of 500-550 C. 
The reaction is exothermic, so that after starting it proceeds 
automatically. An alloy of iron and bismuth or iron and 
copper has occasionally been substituted for the platinum ; 
different catalysts require different temperatures for maxi- 
mum efficiency. 

Properties. Nitric acid is a colourless fuming liquid of 
sp. gr. 1-53 and m.p. 41. Ordinary " commercial " con- 
centrated nitric acid contains 68 per cent, by weight of pure 
acid. When nitric acid is boiled it partially dissociates into 
nitrogen peroxide, oxygen, and water 

4HN0 8 ;= 4N0 2 + 2 + 2H 2 0. 

Nitric acid is the strongest acid known, being equally strong 
with hydrochloric acid. It is also a powerful oxidizing agent, 
and these two characteristics often clash. When nitric acid 
acts upon metals, we may imagine that hydrogen is first 


formed, in the usual way, but that the excess of nitric acid 
then oxidizes the hydrogen and is itself reduced. Reduction 
of nitric acid may go on in various stages 
(i) 2HN0 8 + 2H = 2H 2 + 2N0 2 , nitrogen peroxide. 

(ii) 2HN0 3 + 4H = 3H 2 + N 2 3 , nitrous anhydride. 

(iii) 2HN0 3 + 6H = 4H 2 + 2NO, nitric oxide. 

(iv) 2HN0 8 + 8H = 5H 2 + N 2 0, nitrous oxide. 

(v) 2HN0 8 + 10H = 6H 2 O + N 2 , nitrogen. 

(vi) 2HN0 3 + 12H = 4H 2 + 2NH 2 OH, hydroxylamine. 
(vii) 2HN0 3 + 16H = 6H 2 O + 2NH 3 , ammonia. 

Any or all of these products may therefore be obtained when 
nitric acid acts upon a metal. Of course, in the cases of the 
basic substances ammonium and hydroxylamine, the nitrates 
of these bases would be obtained, and not the free bases. 
With magnesium, dilute nitric acid reacts so rapidly that some 
of the hydrogen escapes oxidation and may be detected in 
the gases evolved. 

The stage to which the reduction of nitric acid is carried, in 
the reaction of this substance with metals, depends upon 
various factors 

(i) Concentration of the acid. Since solutions of metallic 
nitrates have very little oxidizing power, we can conclude that 
the oxidizing powers of nitric acid are due chiefly to the 
undissociated HN0 3 molecules and not to the N0 3 7 ions. 
But nitric acid is a strong acid, i.e. it readily dissociates in 
solution. We should therefore expect that dilution of the 
acid would cause a great diminution in oxidizing power, and 
this is indeed the case. 

(ii) The temperature. Reduction as a rule proceeds further 
with rise of temperature, but this is by no means always true. 

(iii) The nature of the metal. Powerful reducing agents, 
such as zinc, will carry the reduction to a lower stage than 
will metals like copper. 

(iv) The nature of the products of the reaction, if these remain 
in solution. 

(i) Lead. Pb + 4HN0 8 = Pb(N0 8 ) 2 + 2H a O + 2NO t . 


(u) Copper. 3Cu + 8HN0 8 = 3Cu(N0 8 ) a + 4H 2 + 2NO. 
(iii) Zinc. 4Zn+10HN0 3 = 4Zn(N0 3 ) 2 + 3H 2 + NH 4 N0 8 . 
(iv) Tin. 5Sn + 20HN0 8 = H 10 Sn 6 15 + 5H 2 + 20N0 2 . 

H 10 Sn 6 15 is a white insoluble substance called metastannic 
acid ; on heating it decomposes into stannic oxide and water 

H 10 Sn 5 15 = 5Sn0 2 + 5H 2 0. 

It must be realized that these equations are only approxi- 
mate, since the actual reactions are much more complicated 
and vary according to the conditions. 

(v) Sulphur > sulphuric acid, H 2 S0 4 . 
(vi) Phosphorus > phosphoric acid, H 3 P0 4 . 
(vii) Iodine > iodic acid, HIO 3 . 
(viii) Arsenic > arsenic acid, HAs0 8 . 

(ix) Ferrous sulphate reduces nitric acid to nitric oxide, 
which combines with excess of ferrous sulphate to form a 
dark brown compound (FeS0 4 ) 2 .NO. This is the principle of 
the " brown-ring " test for nitrates. The suspected nitrate is 
dissolved in a little water and a few drops of ferrous sulphate 
solution added. Sulphuric acid is then carefully poured in 
and sinks to the bottom as a lower layer. Where the two 
layers meet, a brown ring is formed if a nitrate is present. 
Nitric acid is first liberated by the sulphuric acid and then 
reduced by the ferrous sulphate as above 

6FeS0 4 + 3H 2 S0 4 + 2HN0 8 = 3Fe 2 (S0 4 ) 3 + 2NO + 4H a O 

Ferric sulphate. 

4FeS0 4 + 2NO = 2(FeS0 4 ) 2 NO. 

(x) With bases, nitric acid yields nitrates, all of which 
are soluble with the exception of that of an organic base, 
diphenylendoanilodihydrotriazole. This substance is there- 
fore used, under the name of " Nitron," as a test for nitric 
acid or a nitrate, with a solution of which it gives a white 
precipitate of " Nitron " nitrate. 
Action of heat on the nitrates 

A. Potassium and sodium nitrates first melt, then evolve 
oxygen ; the nitrite of the metal is left 

2KNO, = 2KNO a + O,. 


B. Ammonium nitrate yields nitrous oxide 

NH 4 N0 3 = N 2 + 2H 2 0. 

C. Nitrates of heavy metals, except silver and mercury, 

yield the oxide of the metal, nitrogen peroxide, and 

2Pb(N0 3 ) 2 == 2PbO + 4N0 2 + 2 . 

D. Silver and mercury nitrates yield the metal, nitrogen 

peroxide, and oxygen 

Hg(N0 3 ), = Hg + 2N0 2 + 2 . 

Solid nitrates are powerful oxidizing agents and are there- 
fore used in the preparation of explosives. Gunpowder, for 
instance, consists of potassium nitrate (75 per cent.), with 
sulphur (10 per cent.), and charcoal (15 per cent.). 

(xi) With organic compounds, nitric acid acts in two ways : 
(a) as an oxidizing agent, or (b) as a nitrating agent. 

(a) Dilute nitric acid will readily oxidize compounds of the 
hydrocarbon benzene, C C H 6 . Thus, toluene, C 6 H 5 .CH 3 , is 
oxidized by dilute nitric acid to benzoic acid, C 6 H 5 . COOH. 
Xylene, C 6 H 4 (CH,)j, is converted into phikalic acid, 
C,H 4 (COOH) 2 . 

(ft) Upon the " fatty " compounds, of which methane, 
CH 4 , is the parent substance, dilute nitric acid has little action, 
but when it does react it usually introduces one or more 
NO 2 groups into the molecule ; this action is called nitration. 

(y) Concentrated nitric acid nitrates benzene compounds. 
Thus if benzene itself is shaken with a mixture of nitric and 
sulphuric acids, nitrobenzene, C 6 H 5 N0 2 , is formed in the cold, 
and tfinitrobenzene, C fl H 4 (N0 2 ) 2 , on heating 

C 6 H 6 + HN0 3 = C 6 H 5 .N0 2 + H 2 0. 
C 6 H 9 + 2HN0 8 = C 6 H 4 (N0 2 ) 2 + 2H 2 0. 

These nitrocompounds are often highly explosive, e.g. 
lyddite or picric acid, C 6 H 2 (N0 2 ) 3 .OH, made by acting 
upon phenol, C 6 H 5 OH, with nitric and sulphuric acids, and 
T.N.T. or trinitrotoluene, C 6 H 2 (NO) 3 .CH 3 , were both largely 
ased in the war as high explosives. They burn quietly if 


ignited, but explode violently if detonated by fche explosion 
of a " cap " of mercury fulminate or lead hydrazoate, etc. 

(6) Concentrated nitric acid on fatty compounds acts as an 
oxidizing agent. Thus, if cane-sugar, C 12 H 22 11 , is warmed 
with the concentrated acid, volumes of brown fumes (nitrogen 


peroxide) are given off, and a solution of oxalic acid, | , 

an oxidation product of sugar, is left. 

(e) With certain organic compounds containing hydroxyl 
groups ( OH), such as glycerol and cellulose, nitric acid 
forms true nitrates, for instance, 

CH 2 . OH HN0 8 CH 2 . N0 3 

I I 

CH . OH + HN0 3 = CH . N0 3 + 3H 2 0. 

I I 

CH 2 . OH HN0 3 CH 2 . N0 3 

Glycerol. Glyceryl trinitrate. 

Glyceryl trinitrate is a colourless very explosive oil com- 
monly but incorrectly called " nitroglycerine." The above 
reaction is reversible ; hence sulphuric acid is added to absorb 
the water as soon as it is formed and thus prevent the reverse 
action. Dynamite is a powder made by absorbing " nitro- 
glycerine " in kieselguhr, a siliceous earth composed of the 
remains of minute sea-organisms. In America, it is usual to 
replace the kieselguhr by wood-powder. 

The action of corcentrated nitric acid upon cellulose 
(cotton- wool) results in the formation of the so-called " nitro- 
celluloses " or gun-cottons and collodions. Gun-cotton has the 
empirical formula C 12 H 14 (N0 3 ) 6 4 ; it is used as a " brisant " 
explosive, but not as a " propellant." It has to be detonated ; 
it burns quite quietly if ignited. 


Nitrous oxide, N 2 0. This substance was discovered by 
PRIESTLEY in 1772, but more carefully investigated by DAVY 
(1799). Davy showed that it could be made by heating 
ammonium nitrate 


NH 4 N0 3 = N 2 O + 2H 2 O, 

and called it nitrous oxide. 

Davy's method still remains the best way of preparing 
nitrous oxide, but as the decomposition of ammonium nitrate 
occasionally becomes explosively violent, it is safer to use a 
mixture of sodium nitrate and ammonium sulphate. 

Nitrous oxide is also formed when nitric acid is reduced by 
means of stannous chloride and hydrochloric acid 

2HN0 3 + 4SnCl 2 + 8HC1 = 4SnCl 4 + 51I 2 + N 2 0. 

The gas may be purified from nitric oxide (which is often 
present as an impurity in nitrous oxide made from ammo 1 - 
nium nitrate) by passing through ferrous sulphate solution, 
when the nitric oxide is retained as the brown compound 
(FeSO 4 ) 2 NO. 

Properties. Nitrous oxide is a colourless gas with a sweet 
and not unpleasant smell. It is soluble in cold water and is 
therefore usually collected over hot water, although, since it 
has a density of 22 (air = 14-4), it may conveniently be 
collected by downward displacement when cold. If the gas 
is required dry, it is passed through strong sulphuric acid 
and collected over mercury. 

Nitrous oxide is used as an anaesthetic in dentistry and 
minor surgical operations. It is sometimes called " laughing- 
gas/' since the inhalation of a mixture of nitrous oxide and 
air produces hysterical laughter. " The celebrated Mr. 
Wedgewood, after breathing the gas for some time, threw the 
bag from him, and kept breathing on laboriously with an open 
mouth, holding his nose with his fingers, without power to 
remove them, though aware of the ludicrousness of his situa- 
tion ; he had a violent inclination to jump over the chairs 
and tables, and seemed so light, that he thought he was 
going to fly." Nitrous oxide is an endothermic compound, 
and as it is very readily decomposed on heating into a mix- 
ture of nitrogen and oxygen, containing 33-3 per cent, by 
volume of the latter, substances burn very readily in it 
nearly as well as in pure oxygen. The heat given out during 
the decomposition of the nitrous oxide assists in raising the 



temperature and therefore aids the combustion. A glowing 
splint is re-lit by nitrous oxide, but the latter gas may easily 
be distinguished from oxygen by the following tests 

(i) Nitrous oxide has a characteristic sweet smell. 

(ii) Nitrous oxide is much more soluble in cold water than 
is oxygen. 

(iii) Nitrous oxide when mixed with nitric oxide gives no 
brown fumes of nitrogen peroxide, such as are given by 
oxygen with nitric oxide. 

Fio. 78. Preparation of Nitric Oxide. 

(iv) On burning a piece of phosphorus in nitrous oxide, the 
residual gas on cooling has the same volume as that of the nitrous 
oxide started with 

5N a 

2P - P 2 5 

5N 2 . 

The last experiment shows us that 1 volume of nitrous 
oxide contains 1 volume of nitrogen, 

.'. by Avogadro's Hypothesis, 

1 molecule of nitrous oxide contains 1 molecule of nitrogen. 
/. formula is NgO^.. 

9 is found by a vapour density determination, 
VJD. = 22 /. M.W. = 44. 


Of these 44 parts, 28 are nitrogen, .'. 16 are oxygen. But 16 
is the A.W. of oxygen .". x = J and the formula is N 2 0. 

Nitric oxide, NO. This oxide of nitrogen must of course 
have been obtained by the alchemists, two of the commonest 
of whose chemicals were nitric acid and copper, but it was 
first isolated and investigated by PRIESTLEY in 1772. 
Priestley obtained it by acting upon mercury or copper with 
nitric acid, and called it nitrous air. It is still usually pre- 
pared by Priestley 's method 

3Cu + 8HN0 3 = 3Cu(N0 3 ) 2 + 4H 2 + 2NO. 

The gas obtained in this way, however, is impure. It may 
be purified by absorption in ferrous sulphate solution with 
formation of the brown compound (FeS0 4 ) 2 NO, followed by 
heat, when the brown compound splits up, yielding pure 
m'tric oxide. 

A pure gas may be obtained directly (i) by heating potassium 
nitrate and ferrous sulphate with dilute sulphuric acid, or (ii) 
by warming mercury with a solution of sodium nitrate in 
strong sulphuric acid 
(i) 2KN0 8 + 5H 2 S0 4 + 6FeS0 4 = 

3Fe 2 (S0 4 ) a + 2NO + 4H 2 + 2KHS0 4 . 
(ii) 6Hg + 2HN0 3 + 3H 2 S0 4 - 3Hg 2 S0 4 + 4H 2 + 2NO. 

Properties. Nitric oxide is a colourless gas, which immedi- 
ately combines with free oxygen when the two are brought 
into contact ; it is therefore impossible to say whether it has 
any taste or smell. It is an endothermic substance, but more 
stable than nitrous oxide ; it therefore does not support 
combustion unless the temperature of the burning substance 
introduced is sufficiently high to bring about rapid decom- 
position' of the surrounding nitric oxide. If this decomposi- 
tion occurs, the burning substance goes on burning in the 
oxygen liberated. Feebly burning phosphorus is extin- 
guished if plunged into nitric oxide, but if strongly burning it 
continues to burn in the gas. 

The composition of nitric oxide is determined by enclosing 
a measured volume in a eudiometer tube over mercury. The 
eudiometer contains a spiral of thick iron wire which may be 


raised to a red heat by means of an electric. current. On 
heating, the wire reacts with the nitric oxide, forming solid 
oxide of iron (whose volume is practically equal to that of 
the iron from which it came, and may therefore be neglected) 
and nitrogen. After cooling, it will be found that the volume 
of the residual nitrogen is half that of the nitric oxide taken. 

1 volume of nitric oxide contains volume of nitrogen, 
/. by Avogadro's Hypothesis, 

1 molecule of nitric oxide contains \ molecule of nitrogen, 
/. formula is NO^. 

x is found as for nitrous oxide. 

Note that the composition of nitric oxide cannot con- 
veniently be determined by explosion with hydrogen, as these 
two gases do not form an explosive mixture, unless a trace 
of nitrous oxide is present. 

Nitric oxide is practically insoluble in water and may 
therefore be collected at the pneumatic trough ; it is soluble 
in ferrous sulphate solution, nitric acid, acidified potassium 
permanganate and other substances. A mixture of carbon 
disulphide vapour and nitric oxide burns with a bright bluish 
flame, formerly used by photographers for taking flashlight 

Nitrogen trioxide, N 2 3 , or nitrous anhydride, is formed 
by the action of nitric acid upon arsenious oxide 

2H 2 + As 2 3 + 2HN0 3 = 2H 3 As0 4 + N 2 8 . 

Arsenic acid. 

The gas is dried by means of phosphorus pentoxide and 
then passed through a U-tube surrounded by a freezing- 
mixture, when a deep blue liquid is obtained. If this is 
further dried by standing over phosphorus pentoxide for 
some weeks, it may be boiled without decomposition, and the 
vapour density shows that the molecules are N 4 6 . In the 
presence of even a trace of moisture, decomposition begins, 
first into the monomolecular form, N a 8 , and then into 
N0 2 + NO. 

Ordinary gaseous " N 2 3 " is almost completely dis- 
sociated into nitrogen peroxide and nitric oxide. This mix- 



ture is sometimes called " nitrous fumes " ; on solution in 
ice-cold water it yields nitrous acid 

N 2 3 + H 2 = 2HN0 2 , 
and in an alkali, a nitrite 

N 2 3 + 2NaOH = 2NaN0 2 + H 2 O. 

Nitrogen trioxide is therefore nitrous anhydride, the anhy- 
dride of nitrous acid. 

Nitrogen peroxide, N0 2 or N 2 4 , or nitrogen dioxide, or 
nitrogen tetroxide. This is the commonest oxide of nitrogen. 

FIG. 79. Preparation of Nitrogen Peroxide. 

It is formed when nitric oxide comes into contact with free 
oxygen, when nitric acid acts upon certain metals and various 
organic compounds, and when the nitrates of heavy metals 
are heated. It is generally prepared in small quantity by 
the action of heat upon lead nitrate 

2Pb(N0 3 ) 2 = 2PbO + 4NO a + 2 . 

The mixture of gases is passed through a U-tube surrounded 
by a freezing-mixture ; the oxygen passes on while the nitro- 
gen peroxide condenses to a yellowish liquid. 

A better method of obtaining nitrogen peroxide on a larger 


scale is to pass sulphur dioxide into fuming nitric % acid, keeping 
the liquid cold, until no more will dissolve. On heating the 
product, nitrososulphuric acid, with sodium nitrate a steady 
stream of fairly pure nitrogen peroxide is evolved 

(i) S0 2 + HN0 3 = HS0 4 .NO. 
(ii) HS0 4 .NO + NaN0 3 = N 2 4 + NaHSO 4 . 

Properties. Nitrogen peroxide is at ordinary temperatures 
a reddish brown gas, consisting of a mixture of NO 2 and N 2 4 
molecules. It is convenient to give the name " nitrogen 
dioxide " to the gas consisting entirely of N0 2 molecules, 
" nitrogen tetroxide " to the liquid and solid forms consisting 
entirely of N 2 4 molecules, and to call the ordinary gas 
" nitrogen peroxide." The dissociation of this substance has 
already been discussed (p. 83). 

Nitrogen tetroxide is a pale yellow crystalline solid melting 
at 11 to a yellow liquid. This liquid boils under atmo- 
spheric pressure at 22. At 11 it consists practically 
completely of N 2 4 molecules, but as the temperature rises 
the liquid darkens and partially dissociates into N0 2 mole- 
cules. At the boiling point it consists of about 86 per cent. 
N 2 N 4 molecules and 14 per cent. N0 2 molecules. At 150, 
the gas consists entirely 01 nitrogen dioxide molecules, N0 2 . 
Further heating causes dissociation into nitric oxide and 

N 2 4 ^r N 2 4 ^ 2N0 2 ;r 2NO + 2 . 

Pale yellow Pale yellow Red gas. Colourless gas. 
solid. liquid. 

Nitrogen peroxide dissolves in water to give a mixture of 
nitrous and nitric acids 

N 2 4 + H 2 = HN0 2 + HN0 8 , 
and in alkalis to give a mixture of nitrate and nitrite 

N 2 4 + 2NaOH == NaN0 3 + NaN0 2 + H 2 O. 
It is therefore a " mixed anhydride," being half-way between 
nitrous anhydride, N 2 3 , and nitric anhydride, N 2 5 . 

It is a very strong oxidizing agent. It does not support 
combustion unless the temperature of the burning substance 


is high enough to decompose the gas into nitrogen and oxygen. 
Thus a splint or taper is extinguished, but strongly burning 
phosphorus continues to burn. A mixture of hydrogen and 
nitrogen peroxide if passed over platinum black yields 
ammonia and steam. Carbon monoxide is oxidized fco carbon 
dioxide by nitrogen peroxide, even at ordinary temperatures. 

Interesting compounds called nitroxyls have been obtained 
by SABATIER and SENDERENS by the action of nitrogen 
peroxide on finely divided metals in the cold. They have the 
formula M(N0 2 ) a; , where M is an atom of the metal. (Cf. the 
carbonyls, p. 301.) 

Nitrogen peroxide converts moist sulphur dioxide into 
sulphuric acid 

S0 2 + H 2 + N0 2 = H 2 S0 4 + NO, 

and is used for this purpose in the manufacture of sulphuric 
acid by the lead chamber process (p. 469). 

Nitrogen pentoxide, N 2 S , or nitric anhydride, was dis- 
covered by DEVILLE in 1849, who passed a stream of dry 
chlorine over dry silver nitrate in the cold 

4AgN0 3 + 2C1 2 = 4AgCl + 2N 2 6 + 2 . 

It may also be obtained by mixing phosphorus pentoxide 
with concentrated nitric acid, and distilling off the nitrogen 
pentoxide formed 

2HN0 8 + P 2 6 = 2HP0 3 + N 2 8 . 


Properties. It is a white deliquescent substance which 
explodes on heating and will dissolve in water with evolution 
of heat,, forming nitric acid 

N 2 5 + H 2 == 2HN0 8 . 

It is therefore nitric anhydride. Nitrogen pentoxide is a 
etrong oxidizing agent. 

Nitrous acid, HN0 2 . When nitrogen trioxide is dis- 
solved in ice-cold water nitrous acid is formed 


Salts of this acid with the alkali metals may be obtained 
by heating the corresponding nitrates 

2NaN0 3 = 2NaN0 2 +0 2 

or by passing nitrous fumes (p. 365) into a solution of the 
metallic hydroxide. The acid itself has never been isolated, 
as it very readily decomposes into nitric oxide, nitrogen 
peroxide and water, or into nitric oxide, nitric acid and 

(i) 3HN0 2 = HN0 8 + 2NO + H 2 O. 
(ii) 2HN0 2 ;== NO + N0 2 + H 2 O. 

Potassium and sodium nitrites are the most important 
compounds of nitrous acid. They are pale yellow crystalline 
solids, readily soluble in water. When acted upon by a dilute 
acid they yield nitrous acid, which immediately begins to 
decompose according to the equations given above. The 
acid and its salts may act either as reducing agents or as 
oxidizing agents. 

(a) Oxidizing Reactions. In these reactions the nitrous 
acid may be considered to split up as follows 

4HN0 2 = 4NO + 3 + 2H 2 0. 
Thus, with potassium iodide free iodine is obtained 

2KI + 2HN0 2 = 2KOH + I 2 + 2ND ; 
sulphur dioxide is oxidized to sulphuric acid 
S0 2 + 2HN0 2 = H 2 S0 4 + 2NO ; 

and stannous chloride in the presence of hydrochloric acid ia 
oxidized to stannic chloride 

SnCl 2 + 2HC1 + 2HN0 2 = SnCl 4 + 2H 2 + 2NO. 

(b) Reducing Reactions. In these changes, the nitrous acid 
is oxidized to nitric acid and can therefore bring about reduc- 

2HN0 2 + 2 = 2HN0 8 . 

Thus nitrous acid reduces potassium permanganate solution 
acidified with sulphuric acid, especially on warming 

2KMn0 4 4- 3H,S0 4 + 5HN0 2 

= 5HNO, + KjSO* + 2MnS0 4 + 3H 2 O. 


This reaction is used as a means of estimating the strength 
of a solution of nitrous acid or a nitrite. It is best to add a 
measured volume of the solution of the nitrite to a measured 
volume, in excesa, of standard permanganate, acidified as 
usual ; the liquid is then warmed and the excess of perman- 
ganate estimated by titration with standard ferrous ammonium 
sulphate or oxalic acid. 

Nitrites occasionally occur in drinking-water, where their 
presence probably indicates contamination of the water by 
sewage or other organic matter. They are tested for in water 
analysis by means of a solution of w-phenylenediamine, 
C 6 H 4 (NH 2 ) 2 , in excess of hydrochloric acid. This reagent 
gives a brown colour with a nitrite. 

Sodium nitrite is largely used in the aniline dye industry. 

Hyponitrous acid, H 2 N 2 2 , is obtained as a white flaky 
crystalline solid by evaporating the ethereal solution made by 
adding silver hyponitrite to a dry solution of hydrogen 
chloride in ether. The silver chloride precipitated in the 
action is filtered off before the solution is evaporated. 
Ag 2 N 2 2 + 2HC1 = 2AgCl + H 2 N 2 O 2 . 

Hyponitrous acid is an explosive substance. The sodium 
salt is prepared by reducing sodium nitrite with sodium 
amalgam and water ; the silver salt may be obtained as a 
yellow precipitate by double decomposition of the sodium salt 
and silver nitrate. 

Nitrohydroxylaminic acid, H 2 N 2 3 . When methyl 
nitrate, CH 3 NO 3 , is added to a solution of hydroxylamine and 
caustic soda in methyl alcohol, a white precipitate of the 
sodium salt of nitrohydroxylaminic acid, Na 2 N 2 3 , is obtained. 
The free acid itself is unknown. Addition of dilute hydro- 
chloric acid to the sodium salt causes a vigorous effervescence 
with liberation of nitric oxide. Nitric oxide may perhaps be 
regarded as the anhydride of nitrohydroxylaminic acid, 
though the reaction 

2NO + H 2 = H 2 N 2 0, 
has never been observed. 

Constitution of Nitric and Nitrous Acids. Nitric acid 


is considered to have the structure represented by the for- 


mula yN H. Nitrous acid could then be either 



>N H or - N H. 

In practice, it is found that sodium nitrite behaves in two 

different ways, as if it had both the structure ^N Na, 


and also that expressed by the formula = N ONa, It 
is an example of a tautomeric substance. (Cf. sulphurous 
acid, p. 466.) 

The Nitrogen Cycle. Nitrogen is an essential constituent 
of living matter. Plants require it to be presented to them in 
the form of nitrates, which they take up in solution through 
their roots, from the soil. In the course of the life-process of 
the plant, the nitrogen of the nitrates becomes converted into 
complex nitrogenous organic compounds called proteins. The 
fate of these proteins is twofold. The plant may die, in which 
case the proteins it contains are returned to the soil ; or it 
may be eaten by an animal, when the protein is digested and 
becomes a part of the animal organism. During the life of 
the aiu'mal, proteins are continually being used up and the 
nitrogen excreted in the form of a less complicated organic 
substance such as urea (man) or hippuric acid (horses, etc.). 
If, therefore, dead plants and the excreta and dead bodies of 
animals are returned to the soil, the latter will not become 
impoverished of nitrogen, but the nitrogen returned in this 
way will be of no direct use to plants, which must have it in 
the form of nitrates. Luckily there is an agency at work 
which converts organic nitrogenous compounds in the soil into 
nitrates, This agency is the bacterial flora. The soil swarms 
with bacteria, which cause " decay " of nitrogenous organic 
matter. The first product is ammonia ; by the aid of the 
oxygen of the air one class of bacteria (the nitrite- forming 



bacteria) converts the nitrogen of this ammonia into nitrites, 
and a second kind converts the nitrites into nitrates. (These 
two classes of bacteria are called the nitrifying bacteria.) In 
this way the cycle of changes undergone by the nitrogen in 
nature is complete (Fig. 80). 

There are, however, several factors that modify this 
simple " nitrogen cycle." In the first place, there are certain 
species of soil bacteria that convert nitrates into nitrogen, 
which escapes into the air and is therefore lost ; these are the 


Atmospheric Nitrogen 

>> Organic Ifilrogen 

Fio. 80. 

denitrifying bacteria. Secondly, a great deal of nitrogen is 
taken from the soil and never replaced, owing to our wasteful 
system of turning sewage into the sea whenever possible. 
Thirdly, in order to get sufficiently large crops to feed the 
population of the world, large quantities of nitrates or ammo- 
nium salts have to be added to so 1, as artificial manures. 
Fourthly, certain plants, such as peas beans, and clover, have 
the power of making direct use of atmospheric nitrogen, by 
means of peculiar bacteria which live in nodules on their roots ; 
but the enrichment of the soil brought about in this way and 
by the production of nitric acid in the air during thunder- 


storms is not sufficient to make up for the annual drain on 
nitrogen content of the soil caused by growing and removing 
a crop. 

In other words, for the establishment of a true equilibrium 
in the nitrogen cycle, at present much nitrogen has to be added 
to the soil, by man, in the form of nitrates or ammonium 
salts. The production of ammonium salts in the manufacture 
of coal-gas is far too little to supply the need, and the beds of 
Chile saltpetre are being rapidly exhausted. Hence the great 
importance of the recently-perfected methods of synthesizing 
nitric acid and ammonia, a problem to which the attention of 
chemists was fortunately directed in good time by the fore- 
sight of the late SIR WILLIAM CROOKES. 

The complete Nitrogen Cycle may be represented by Fig. 
80 on the previous page. 


1. Describe the synthetic production of ammonia. 

2. What are the chief properties of ammonia ? What evidence 
have we that in ammonium compounds the nitrogen is pentavatent T 

3. Compare the properties of ammonium salts with those of the 
alkali metals. 

4. Describe the preparation and properties of hydroxylamine, 
hydrazine, and hydrazoic acid. 

5. What oxides does nitrogen form ? How are they made and 
what are their chief properties ? 

6. What do you understand by the Nitrogen Cj^clo ? 

7. Write an account of the various methods employed for the fixation 
of atmospheric nitrogen. 

8. Discuss the action of nitric acid upon (a) metals, (6) organic 

9. Are nitrites oxidizing agents or reducing agents ? Illustrate your 
answer with examples. How would you estimate the strength of a 
solution of sodium nitrite ? 


Group in Periodic System : V ; Symbol : P ; Valency : 3 
and 5; Atomicity of Vapour: 4; Atomic Weight: 31-02; 
Melting Point : 44 ; Specific Gravity : 1-83. 

History. During the course of his chemical researches^ 
BRAND of Hamburg discovered phosphorus. He made it, in 


1669, by evaporating large quantities of urine to a thick 
syrup, mixing this with sand, and distilling. Brand sold the 
secret of his discovery to KRAFFT, who made a grand tour of 
Europe exhibiting the marvel at various Courts, including 
that of Charles II. The striking properties of the new sub- 
stance excited the curiosity of chemists, and KUNCKEL (1678) 
and BOYLE (1680) independently discovered the way to pre- 
pare it. It was for long called English phosphorus, apparently 
on account of its preparation by Boyle. 

In 1771 GAHN showed that phosphorus is present in bones, 
and SCHEELE devised a means of extracting it from bone-ash. 1 
In 1772 LAVOISIER concluded that phosphorus was an element. 

Occurrence. Phosphorus is too reactive an element to 
exist naturally in the uncombined state. It is widely distri- 
buted, though in comparatively small quantity, in the form of 
phosphates or salts of phosphoric acid, H 3 PO 4 . The chief 
naturally occurring phosphates are phosphorite and sombrerite, 
impure Ca 3 (P0 4 ) 2 ; apatite, 3Ca 3 (P0 4 ) 2 .CaF 2 ; chlorapatfye, 
3Ca 3 (P0 4 ) 2 .CaCl 2 ; vivianite, Fe 3 (P0 4 ) 2 .8H 2 ; wavellite, 
[A1 2 (P0 4 ) 2 ] 2 .A1 2 (OH) 6 .9H 2 ; Redonda phosphate (impure 
calcium phosphate) ; and coprolites (fossilized excreta contain- 
ing calcium phosphate). It is an essential constituent of aD 
living matter, and its presence in soil is therefore of the great- 
est importance. Soils poor in phosphorus have to be treated 
with phosphatic manures, such as basic slag (p. 549), or calcium 
superphosphate (p. 248). Calcium phosphate is the chief 
constituent of bones. 

Preparation. The actual details of modern manufacturing 
processes for the preparation of phosphorus are jealously 
guarded trade secrets, but the general principles are known. 

(i) Extraction from Bones. The bones are first treated with 
organic solvents such as benzene, carbon disulphide, ether, 
or acetylene tetrachloride (which has the great advantage of 
being non-inflammable) to remove fats ; the residue is then 
heated with water under pressure, when a solution of glue is 
obtained. The solid left is then destructively distilled in iron 

1 A boy onoe told me in reply to a- question on phosphorus, 
" phosphorus was made by Scheele from Beau Nash " 1 



retorts in absence of air, when a distillate called bone- oil or 
DIPPEL'S oil comes over, and a mixture of finely divided carbon 
and calcium phosphate is left. This mixture is called animal 
charcoal, or bone black. It is employed as a decolourizing 
agent in the refinement of sugar, and after use in this way is 
burnt in the air, the residue consisting of bone-ash (about 
80-85 per cent, calcium phosphate). 

The bone- ash is treated with an equivalent weight 
of concentrated sulphuric acid, when calcium sulphate 

and a solution of 
phosphoric acid are 

Ca 3 (P0 4 ) 2 +3H 2 S0 4 
=3CaS0 4 + 2H 3 P0 4 

Orthophosphoric acid. 

The calcium sul- 
phate, which is in- 
soluble, is allowed to 
settle and the clear 
solution of phosphoric 
acid run off and 
evaporated till it is 
of a syrupy con- 
sistency. The 
acid, HP0 3 

FIG. 81. Bone Ash Phosphorus Furnace. 
A. Retorts. B. Distilling pipe. D. Collecting pots 

syrup consists of mefophosphoric 

H 3 P0 4 = H 2 + HP0 3 . 

It is mixed with about one quarter of its weight of crushed 
and powdered coke and the mixture strongly heated in retorts 
of Stourbridge clay 

2HP0 8 + 6C = H 2 + 6CO + 2P. 

Phosphorus vapour distils off and is condensed under water 
to a yellow liquid, which is purified by filtration through 
chamois leather and in other ways not known to the public. 
On cooling, the liquid phosphorus sets to a yellow crystalline 
solid. It is usually cast in sticks. 

The extraction of phosphorus from bones is now obsolescent, 


and will doubtless soon be replaced entirely by the modern 
process described in the next paragraph. 

(ii) Extraction from Mineral Phosphates. Nearly 100 years 
ago the German chemist WOIILBB suggested that phosphorus 
could be made by heating mineral phosphates with sand and 
coke to a very high temperature. This method was not used 
until 1898 owing to the difficulty of procuring at a reasonable 
cost the high temperature necessary for the reaction to occur, 
but in 1898 RE ADMAN, ROBINSON, and PARKER solved the 
problem by using an electric furnace. Since that date, most 
of the phosphorus of commerce has been made by the electri- 
cal process, by the firm above-mentioned, by ALBRIGHT 
and WILSON (Niagara Falls), and by VIOLET (Paris and 

The phosphates are mixed with sand, Si0 2 , and coke in the 
proportion (determined by analysis of the ore) to correspond 
with the equation 

Ca 3 (P0 4 ) a + 50 + 3Si0 2 = 3CaSi0 3 + 2P + 5CO. 

Calcium silicate. 

The mixture is then heated to a high temperature by means 
of the electric arc struck between two stout carbon rods in a 
strong brickwork furnace. The calcium silicate melts and 
sinks to the bottom, whence it is run off as required ; phos- 
phorus vapour and carbon monoxide escape through a pipe at 
the top, and fresh mixture is added continuously by means of 
a hopper and screw arrangement, as shown in the figure. The 
process is thus continuous. 

The phosphorus vapour is condensed under water in copper 
vessels and is purified as described in (i), or by melting under a 
solution of potassium dichromate and dilute sulphuric acid, 
which oxidizes the impurities and removes them from the 
phosphorus, or by redistillation, etc. 

Properties. When freshly prepared, phosphorus is a 
translucent pale yellow waxy crystalline solid, which can 
easily be cut with a knife. It gradually darkens on exposure 
to light, becoming finally very dark brown. 

Its specific gravity is 1-8. It melts at 44 and boils at 290. 



The liquid and vapour are colourless, and vapour density 
determinations show that the molecule of the vupour is P. 
At high temperatures partial dissociation occurs 

P 4 

P 2 + P 2 . 

Phosphorus is practically insoluble in water but readily 
dissolves in many organic solvents, such as carbon disulphide, 

and also in phosphorus 
trichloride. It ignites in 
moist air at 30 C. 

The name " phosphorus " 
(light-bearer) was given to 
the element on account of 
its peculiar property of glow- 
ing in the dark. NICOLAS 
LEMERY, in his Cours de 
Ghimie (1694), mentions 
certain practical jokes which 
had been carried out by 
means of the phosphores- 
cence and ready ignition 
of phosphorus. He says 
that a piece of phosphorus 
was placed in the bed of a 
visitor to Boyle ; a servant 
had to extinguish the con- 
flagration by throwing 
buckets of water over both 
bed and visitor. 
The glow is due to slow oxidation of the phosphorus, chiefly 
to phosphorous oxide, P 4 6 . When phosphorus is burnt in a^r 
or oxygen the main product is phosphoric oxide or phosphorus 
pentoxide, P 2 6 . Phosphorus will not burn in absolutely dry 
oxygen. The glow of phosphorus may be shown in an elegant 
manner by the following experiment. A few pieces of phos- 
phorus are placed in a round-bottom flask with a little water. 
On boiling the water, phosphorus vapour passes up with the 
steam and oxidizes at the mouth of the flask with a greenish 

FIG. 82. Electric Furnace for 
manufacture of Phosphorus. 


flame, which is BO cold that it will not burn paper or even 
set fire to a match. 

Phosphorus is very chemically active. It will combine 
directly with halogens, oxygen and many metals. It is very 
poisonous, the vapour causing a disease called " phossy-jaw," 
or decay of the bones of the jaw (and of other parts of the 
body as well). 

Red Phosphorus. Phosphorus exists in many allotropic 
modifications ; the form already described is called yellow, 
white, or a-phosphorus. If this form is heated in an inert 
atmosphere to 250 (especially in the presence of a trace of 
iodine as catalyst) it is converted into a red modification, red 
phosphorus. This was discovered by SOHKOTTEB in 1845 and 
was considered to be amorphous until 1890, when RETGBKS 
showed that it was minutely crystalline. Red phosphorus is 
important commercially, and is manufactured by heating the 
yellow form in a cast-iron pot to a temperature of 230-250. 
It is important that the temperature should not rise above 
250, as at higher temperatures the reaction may become 
explosive. Thermometers are therefore placed in the pot and 
have to be encased in iron tubes, since hot phosphorus attacks 

Red phosphorus is not so chemically active as the yellow 
variety. It is so insoluble ih water that it is non-poisonous 
when introduced into the alimentary canal ; if, however, it is 
injected into the blood, the characteristic symptoms of phos- 
phorus poisoning make their appearance. It will not dissolve 
in carbon disulphide, and has a high ignition point (260). 
Its specific gravity is 2-25. It does not glow in moist air, and 
will not take fire spontaneously in chlorine, as the yellow form 

Transformation of red phosphorus into yellow, and of yellow 
into red. 

(i) Yellow to red. Heat to 240 in an inert atmosphere, 
with trace of iodine as catalyst. 

(ii) Red to yellow. Heat in an inert atmosphere until the 
vapour is produced (above 550). On rapidly cooling the 
vapour yellow phosphorus is formed. 


There is no point at which yellow and red phosphorus can 
exist in equilibrium together ; the yellow form is always 
unstable with regard to the red. Phosphorus is therefore a 
monotropic substance. (Cf. sulphur, p. 458.) The velocity of 
transformation of yellow into red at ordinary temperatures is 
so small that the yellow form appears to be stable. The 
vapour pressure of yellow phosphorus is greater than that of 
the red form at the same temperature, and the change of the 
yellow into the red is accompanied by evolution of 3,700 
calories per gram-atom. 

Other allotropic forms of phosphorus are known. SCHENCK'S 
scarlet phosphorus is important on account of its use in the 
match industry. It is prepared by boiling a solution of yellow 
phosphorus in phosphorus tribromide, PBr 3 , until no further 
precipitation of the scarlet form occurs. It is more active 
than red phosphorus, but is non-poisonous and does not 
spontaneously oxidize in the air. 

Matches. Matches were invented by CHANCEL in 1805. 
The earliest forms consisted of wooden splints tipped with a 
mixture of potassium chlorate and sugar. To ignite these it 
was necessary to carry a bottle of sulphuric acid into which 
the match could be dipped. In spite of the inconvenience of 
this procedure, Chancel's matches were widely used for nearly 
half a century. Matches which ignite by friction were in- 
vented about 1840 by the Frenchman SAURIA and others, 
including GENERAL CONGREVE, an English officer. They 
consisted of strips of wood headed with a mixture of antimony 
sulphide, potassium chlorate and gum, and were ignited by 
rubbing vigorously on sandpaper. About the same time, a 
mixture of yellow phosphorus, sulphur and potassium chlorate 
was employed in match-making, but the disadvantage of 
these matches was that they often ignited at inconvenient 
moments. Nevertheless, yellow phosphorus matches were 
popular for many years until the effects on the workmen 
(" phossy-jaw ") became so marked that laws were passed 
forbidding the use of such a poisonous and dangerous chemical. 

" Strike anywhere " matches are now tipped with a mixture 
of Schenck's scarlet phosphorus, potassium chlorate, red lead, 


gum and a colouring matter. " Safety " matches are made 
from a mixture of potassium chlorate, antimony sulphide, red 
lead, potassium dichromate and gum. They must be ignited 
by rubbing on a specially prepared surface (on the side of the 
box) containing red phosphorus, powdered sand or glass, 
antimony sulphide and gum. Most of the processes in 
the modern manufacture of matches are carried out by 

COMPOUNDS OF PHOSPHORUS. Phosphorus forms four 

PH 3 , phosphine or " phosphoretted hydrogen." 

P 2 H 4 , liquid hydrogen phosphide. 

P 12 H 6 and P 2 H 2 , solid hydrogen phosphides. 

The true nature of these solids is still doubtful. 

Phosphine, PH 8 . Phosphine was first prepared by 
GENGEMBRE in 1783, by heating yellow phosphorus with 
caustic potash solution. This method is still employed for 
laboratory preparation of the gas 

4P + 3KOH + 3H 2 = 3KH 2 P0 2 + PH 8 . 
KH 2 P0 2 is called potassium hypophosphite. To conduct 
the experiment caustic potash solution and some pieced 
of yellow phosphorus are placed in a round-bottomed flask 
fitted with a cork carrying a delivery tube and another 
tube which admits coal-gas. All air is first swept from the 
apparatus by a stream of coal-gas and the flask is then 
heated. Phosphine comes off, and as each bubble rises from 
the water of the trough into the air it ignites spontaneously 
and forms a vortex ring of white phosphorus pentoxide. 

Phosphine may also be prepared by the action of water 
upon calcium phosphide, Ca 3 P 2 

Ca 3 P 2 +6H 2 = 3Ca(OH) 2 + 2PH 2 ; 
or by heating phosphorous acid 

4H 3 P0 8 = 3HP0 3 + 3H 2 + PH 8 ; 

or by the action of dilute caustic soda solution upon pboa* 
phonium iodide 

PH 4 I + NaOH = Nal + H 2 + PH t . 



The gas prepared in the last two ways is much purer than 
that prepared by the first two, and is not ^spontaneously 
inflammable. The spontaneous ignition of phosphine is 
caused by the presence in the gas of traces of the liquid 
hydrogen phosphide, P 2 H 4 , which is itself spontaneously 
inflammable and thus sets fire to the phosphine. The liquid 
hydride may be removed by passing the gas through a U-tube 
surrounded by a freezing-mixture, which retains the P 2 H 4 but 
allows the phosphine (no longer spontaneously inflammable) 


FIG. 83. Preparation of Phosphine. 

to pass on. Decomposition of phosphonium iodide is the best 
way of getting pure phosphine. 

Phosphine is a colourless gas, only slightly soluble in water 
(contrast ammonia, p. 342). It has an unpleasant smell, 
reminiscent of decayed fish, and is extremely poisonous. The 
presence of traces of calcium phosphide in calcium carbide 
causes phosphine to be present as an impurity in ordinary 
acetylene, hence the bad smell of the gas. Pure acetylene has 
a sweetish and not at all unpleasant odour. Phosphine is a 
basic anhydride, like ammonia, for it will combine directly with 
acids to form phosphonium salts. The base, however, phos- 
phonium hydroxide, PH 4 OH, is unknown, and a solution of 
phosphine in water has no action on litmus. The chief 


phosphonium compound is phosphonium iodide, PH 4 I. It is 
prepared by mixing phosphorus and iodine in a current of 
carbon dioxide and then dropping water slowly on to the 

9P + 51 + 16H 2 = 4H 3 P0 4 + 5PHJ. 

The phosphonium iodide may then be sublimed off by gentle 
heat, and is obtained in the form of beautiful colourless 
crystals ; it is decomposed if added to water or solutions of 
caustic alkalis, with evolution of phosphine. Phosphonium 
iodide and the chloride both dissociate on heating, as does 
ammonium chloride. 

Phosphides of Metals. These substances may be prepared 
by the action of phosphorus or phosphine upon solutions of 
certain metallic salts, such as silver nitrate and copper sul- 
phate, which yield black precipitates of Ag 3 P and Cu 8 P 2 
respectively. They may also be made by passing phosphorus 
vapour over the strongly heated metallic oxide, or by heating 
the phosphate of the metal with carbon in the electric furnace. 
The last method is used in the preparation of calcium phos- 
phide. Commercial calcium phosphide is a reddish brown 
solid that yields spontaneously inflammable phosphine on 
treatment with water. Mixed with calcium carbide it is 
therefore used in Holmes 1 signals, which consist of tins con- 
taining the mixture, attached to a buoy. When required for 
use, the tins are pierced at each end and thrown into the sea. 
The acetylene evolved is ignited by the ignition of the phos- 
phine it contains, and the sea is lit up. 

Halogen Compounds of Phosphorus. Phosphorus com- 
bines with the halogens to form two series of compounds, PX S 
and PX 5 , ^here X = F, Cl, Br, or I. The cluef of these 
compounds are phosphorus trichloride, PC1 8 , phosphorus 
pentachloride, PC1 5 , and phosphorus tribromide, PBr,. In PX, 
the phosphorus atom is tervalent, in PX 6 it is quinquevalent. 

Chlorides. Phosphorus trichloride is prepared by pass- 
ing chlorine through molten phosphorus in a retort from which 
all air has previously been removed by a stream of carbon 
dioxide. The phosphorus takes fire and burns in the chlorine. 



and the phosphorus trichloride distils over as a colourless 
oily liquid. It fumes in the air and is decompbsed by water 
with formation of phosphorous acid and hydrochloric acid 

PC1 3 + 3H 2 O = H 3 P0 3 + 3HC1. 

It is extensively employed in organic chemistry for replace- 
ment of hydroxyl groups by chlorine atoms. 

Phosphorus pentachloride is made by allow- 
ing phosphorus trichloride to react with 
excess of chlorine. A convenient apparatus 
for the purpose is shown in the figure. It 
consists of a wide-necked jar fitted with a 
cork carrying -a dropping-funnel and two 
delivery tubes. A current of dry chlorine 
is slowly passed through the jar, and the 
trichloride run in drop by drop from the 
funnel. Solid phosphorus pentachloride col- 
lects in the jar. It is a yellowish crystal- 
line substance with a peculiar smell. It 
sublimes on heating and the vapour is dis- 
sociated into PC1 3 and C1 2 (pp. 83, 124). 
It fumes in the air and is vigorously attacked 
by water ; the first change results in the 
formation of a colourless oily liquid, phos- 
phorus oxy chloride, POC1 3 , and this is then 

FIG. 84. Prepar- acted upon by more water, yielding orthophos- 


(i) PC1 5 + H 2 = POC1 3 + 2HC1. 

lne chforine'out. ces8 Like the trichloride, phosphorus penta- 

chloride is used in organic chemistry for 

replacing hydroxyl groups by chlorine atoms. Thus, if it is 

added to ethyl alcohol, a violent reaction occurs and ethyl 

chloride is formed 

C 2 H 5 OH + PC1 8 = C 2 H 6 C1 + POC1, + HC1, 

while if benzoic acid and phosphorus pentachloride are ground 
together in a mortar the mass soon liquefies and volumes of 


hydrochloric acid gas coine off, benzoyl chloride being left 

C 6 H 5 .CO.OH + PCi 5 = C 6 H 5 .CO.C1 + P001 3 4- 

Phosphorus tribromide is a colourless liquid, very similar to 
phosphorus trichloride, made by dropping bromine on to red 
phosphorus. The tribromide and tri-iodide are used in many 
organic and inorganic reactions, but are usually made in situ 
when required, by direct action of halogen upon phosphorus. 

Oxides. Phosphorus forms three oxides 1 and a doubtful 
fourth [P 4 ? phosphorus suboxide], phosphorous oxide, 
P 2 3 , phosphorus tetroxide, P 2 4 , and phosphoric oxide (or 
phosphorus pentoxide), P 2 5 . 

Phosphorous oxide, or phosphorus trioxide, P 2 8 , is formed, 
together with the pentoxide, when phosphorus is burnt in a 
limited supply of air. The phosphorus is burnt in a long hard 
glass tube, and the products of combustion passed through a 
plug of glass wool. This stops the pentoxide, which is a 
solid, but the more volatile trioxide, being still in the gaseous 
state, passes through and is condensed in a cooled U-tube. It 
is a white crystalline solid rather waxy in appearance. It 
melts at 23 and boils at 173 ; it rapidly oxidizes in the 
air to the pentoxide, and dissolves in water, slowly in the cold 
but much more rapidly on heating, with formation of phos- 
phorous acid, H 3 P0 3 . It may therefore be called phosphorous 

P 2 3 + 3H 2 = 2H 3 P0 8 . 

The vapour density of the trioxide is 110, corresponding to 
the double formula P 4 6 . As, however, all the chemical 
reactions in which the substance takes part can be satisfac- 
torily expressed by using the formula P 2 3 , it is customary 
to use this formula, and not the double one which is no doubt 
more accurate. Phosphorus trioxide is a poisonous sub- 
stance and smells of garlic. 

Phosphorus tetroxide, P 2 4 , is formed when the trioxide is 
heated under pressure, 4P 2 3 = 3P 2 4 + 2P. 

It is a white crystalline solid, and resembles nitrogen 

1 Besson claims to have obtained P 2 O as well, but this may be mer^ty 
impure red phosphorus. 


peroxide in dissolving in water to give a mixture ot acids ; in 
this case'phosphorous and phosphoric acids. It is therefore 
a " mixed anhydride " 

P 2 4 + 3H 2 = H 8 PO a + H 3 P0 4 . 

Phosphorus pentoxide, P 2 5 , is made by burning phosphorus 
in excess of air or oxygen. The commercial product always 
contains a little trioxide as impurity. The vapour density 
shows that the vapour consists chiefly of P 4 10 molecules, but 
the simple formula P 2 5 is generally used as it satisfactorily 
expresses the chemical behaviour of the compound. 

It is a white crystalline solid with a great affinity for water ; 
it is, indeed, the most effective drying-agent known. It is 
very quickly turned to a semi-liquid mass, metaphosphoric 
acid, HP0 8 , on exposure to air, and dissolves in water with a 
hissing noise forming metaphosphoric acid in the cold but 
orthophosphoric acid if the water is hot 

(i) P 2 5 + H 2 2HP0 8 . 
(ii) P 2 5 + 3H 2 = 2H 8 P0 4 . 

In addition to its drying powers it possesses the property of 
taking the elements of water out of many substances which 
contain them. Thus it yields nitrogen pentoxide with nitric 
acid, sulphur trioxide with sulphuric acid, acetonitrile with 
acetamide, and carbon sub- oxide with malonic acid 

(i) 2HN0 3 - H 2 - N 2 6 . 
(ii) H 2 S0 4 - H 2 = S0 3 . 
(iii) CH 8 .CO.NH 2 - H 2 = CH 3 .CN. 

Acetamide. Acetonitrile. 


(iv) CH 2 2H 2 - 0,0,. 

Malonic acid. Carbon sub-oxide. 

Phosphorus pentoxide is the anhydride of the phosphoric 
acids, orthophosphoric acid y metaphosphoric acid, and pijro- 
phosphoric acid. 

Oxyacids of Phosphorus. Many oxyacids of phosphorus 


are known, and some of them are of importance. We shaD 
consider the following only 

Hypophosphorous acid, H 3 P0 2 . 

Orthophosphorous acid, H 3 P0 8 . 

Hypophosphoric acid, H 2 PO 8 . 

Orthophosphoric acid, H 3 PO 4 . 

Pyrophosphoric acid, H 4 P 2 T . 

Metaphosphoric acid, HPO 3 .. 

Hypophosphorous acid, H 3 P0 2 . A sodium salt of this acid 
is formed when yellow phosphorus is boiled with caustic soda 
solution, as in the preparation of phosphine. Baryta-water 
instead of caustic soda yields the corresponding barium salt 

3Ba(OH) 2 + 8P + 6H 2 O = 3Ba(H 2 PO 2 ) 2 + 2PH 8 . 

Barium hypophosphite. 

If the solution of the barium hypophosphite is first treated 
with carbon dioxide, to remove excess of baryta as carbonate, 
and the filtered solution then evaporated, colourless crystals 
of the salt separate out. 

The free acid may be obtained by addition of the calculated 
weight of sulphuric acid to a solution of the barium salt. 
The precipitated barium sulphate is filtered off and the 
filtrate evaporated. 

Ba(H 2 P0 2 ), + H 2 S0 4 = BaSO 4 = 2H 8 PO 2 . 

As obtained in this way hypophosphorous acid is a colour- 
less syrup which can be frozen to white crystals melting at 
17. It is a weak monobasic acid, and probably has the 
constitution represented by the formula 

O - P^-H, 


the replaceable hydrogen being that in the hydroxyl group. 
Hypophosphorous acid and its salts decompose on heating, 
with evolution of phosphine 

4H 8 PO 2 = 2HP0 8 -f 2H 2 + 2PH 8 . 

The hypophosphites of sodium, potassium and calcium are 
used in medicine, e.g. in " Parrish's Chemical Food.' 1 If 



sodium hypophosphite is added to copper sulphate solution 
a precipitate of copper hydride, CuH, is formed. 

Orthopkosphorous acid, H 3 P0 3 , is formed by dissolving 
phosphorus trioxide in water 

P 2 3 + 3H 2 = 2H 3 P0 3 , 
or by the action of a phosphorus trihalide on water, 

e.g. PC1 3 + 3H 2 = H 3 P0 3 + 3HC1. 

On evaporating the solution until the temperature has risen 
to 180, and then cooling, crystals of orthophosphorous acid 

Phosphorous acid is a white crystalline solid melting at 
72. It is a dibasic acid, and a powerful reducing agent since 
it readily takes up oxygen to go to orthophosphoric acid, 
H 3 P0 4 . Thus it reduces silver, copper and gold salts, in 
solution, to the metals. Its constitution is probably 

= P^ OH, although from its formation from phosphorus 


trichloride and water we should expect it to be P^-OH 


y Cl HOH OH 

P^ Cl + H OH - 3HC1 + P^-OH 

\Cf HOH X)H 

Compare nitrous and sulphurous acids (pp. 370 and 466). 

Organic derivatives of both forms of phosphorous acid 
represented by the above formulae are known. 

On heating, phosphorous acid splits up into phosphine and 
metaphosphoric acid 

4H 3 P0 3 = 3HPO 8 + 3H 2 + PH 8 . 

Hypophosphoric acid, H 2 P0 3 , may be prepared by adding 
yellow phosphorus to silver nitrate solution and decomposing 
the precipitated silver hypophosphate with hydrochloric 

Ag 2 P0 3 + 2HC1 = 2AgCl + H 2 P0 3 . 

It is a white crystalline solid which on heating splits up 


into phosphine and phosphoric acid. It is not a reducing 


agent, and probably has the composition = 

in which the phosphorus is quadrivalent, or the double mole- 


= P< 

may be formed. 

Orthophosphoric acid, H 3 P0 4 , is the most important acid of 
phosphorus. It may be obtained by decomposing calcium 
phosphate with sulphuric acid 

Ca 3 (P0 4 ) 2 + 3H 2 S0 4 = 3CaS0 4 + 2H 3 P0 4 , 
and evaporating the solution after filtration from the calcium 
sulphate. It is also formed when phosphorus pentoxide is 
added to boiling water 

P 2 5 + 3H 2 = 2H 3 P0 4 , 
and when a solution of metaphosphoric acid, HPO 3 , is boiled 

HP0 3 + H 2 = H 3 P0 4 . 

The pure acid may be most conveniently prepared by boiling 
yellow or red phosphorus (preferably the latter) with con- 
centrated nitric acid. The mixture is placed in a round- 
bottom flask fitted with a reflux condenser. 

The aqueous solution of the acid obtained in any of the 
above ways is concentrated until the temperature rises to 
140. The syrupy liquid is then allowed to cool in a desic- 
cator, and the acid separates out as colourless rhombic 
crystals, melting at 41-7. 

Orthophosphoric acid is soluble in water, forming a feebly 
acid solution. It ionizes chiefly into H" and H 2 PO 4 ', although 
even this dissociation is small. The further dissociation 

H- + H 2 P0 4 ' ^= H- + HTO/' 

is practically negligible. Nevertheless, phosphoric acid i* a 
tribasic acid and forms three sodium salts, 


NaH 2 PO 4 , sodium dihydrogen orthophosphate, 
Na 2 HP0 4 , disodium hydrogen orthophoephate, 
Na 3 P0 4 , trisodium orthophosphate, or normal sodium 
orthophosphate . 

Normal sodium phosphate gives a strongly alkaline solution. 

This is easily explained on the ionic theory. Sodium 

phosphate, like all sodium salts, ionizes very largely in 


Na 3 PO 4 ^= 3Na + P0 4 '". 

But orthophosphoric acid is a very weak acid, and the P0 4 '" 
ions immediately react with the hydrogen ions present in 
water to form the H 2 PO/ ion 

H 2 ^z IT + OH' 
2H' + P0 4 '" ^r H 2 PO 4 '. 

The equilibrium between the water molecules and the hydro- 
gen and hydroxyl ions is thus upset, and more water ionizes, 
thus furnishing more hydrogen ions to the PO 4 '" ions. This 
process goes on until practically all of the P0 4 '" ions have 
been converted into H 2 P0 4 ' ions. The hydroxyl ions corre- 
sponding to the hydrogen ions which have been taken up are 
left over and thus the solution has an alkaline reaction. 

The disodium salt, Na 2 HP0 4 , is ordinary laboratory 
" sodium phosphate." It has a slightly alkaline reaction in 

The monosodium salt, NaH 2 P0 4 , has a very slightly acid 
reaction in solution. 

The constitution of orthophosphoric acid is probably 
represented by the following formula 


= Pf-OH. 

Titration with caustic soda using litmus or phenolphthalein 
as indicator gives the disodium salt ; with methyl orange the 
colour change occurs at the stage corresponding to NaH 2 P0 4 . 

Pyrophosphoric acid, H 4 P 2 7 , is formed when ortho- 
phosphoric acid is cautiously heated at a temperature of 


about 215-220. Two molecules of the orthophosphoric acid 
lose one molecule of water 

2H 3 P0 4 = H 4 P A + H 2 0. 

Pyrophosphoric acid is a white crystalline solid melting at 
61. It is a tetrabasic acid, but salts of the types NaH 3 P 2 7 
and Na 3 HP 2 7 are not known only the normal and diacid 
salts of the types Na 4 P 2 7 and Na 2 H 2 P 2 7 have hitherto been 
prepared. If dissolved in water, pyrophosphoric acid changes 
slowly in the cold, but more rapidly on warming, into ortho- 
phosphoric acid. 

Pyrophosphates may often be made by heating the corre- 
sponding orthophosphates ; thus " sodium phosphate/* 
Na 2 HPO 4 , on heating yields sodium pyrophosphate 

2Na 2 HP0 4 = H 2 + Na 4 P 2 O 7 . 

Pyrophosphates in solution gradually revert to ordinary 
sodium phosphate ; the change is accelerated by the addition 
of a little mineral acid. 

The estimation of " P 2 O f " in a solution is often carried 
out by addition of ammonia, ammonium chloride and 
magnesium sulphate, when a precipitate of magnesium 
ammonium orthophosphate is obtained. This is collected and 
heated to redness, when it loses ammonia and water and 
leaves a residue of magnesium pyrophosphate, which is weighed. 

2MgNH 4 P0 4 = H 2 + 2NH 3 + Mg 2 P 2 O 7 . 

Metaphosphoric acid, HP0 3 , may be made by heating the 
ortho or the pyro acid to redness 

H 3 P0 4 = H 2 + HPO, 
H 4 P 2 7 = 2HPO a + H 2 0, 

of by adding phosphorus pentoxide slowly to cold water 
P 2 5 + H 2 = 2HP0 3 . 

It is a glassy transparent solid and is put on the market as 
" glacial phosphoric acid." It is a monobasic acid and the 
chief salt is sodium metaphosphate, NaP0 8 . This is formed 
when sodium ammonium hydrogen orthophosphate or 


microcosmic salt (so-called because it is found in the urine oi 
the " microcosm," i.e. man), is strongly heated 

NaNH 4 HP0 4 = NaP0 3 + H 2 + NH 3 . 

Fused sodium metaphosphate dissolves many metallic oxides 
to form coloured orthophosphates ; microcosmic salt is 
therefore sometimes used instead of borax for the " bead " 
test in analysis. 

Tests for Phosphates. All phosphorus oxyacids and 
their salts when heated with ammonium molybdate solution, 
(NH 4 ) 2 Mo0 4 , and excess of concentrated nitric acid, give a 
yellow precipitate of variable composition called ammonium 
phosphomolybdate. If the precipitation is carried out under 
certain specified conditions, the precipitate has a definite 
composition, and this reaction may then be used for the 
estimation of phosphates or " P 2 5 ." 

With silver nitrate solution, orthophosphates give a yellow 

pyrophosphates give a white 

metaphosphates also give a 

white precipitate. 

With a solution of white of egg, metaphosphates cause coagu- 

ortho and pyro phosphates 
have no action. 

For further tests, see FENTON'S Notes on Qualitative 


Group in Periodic System : V ; Symbol : As ; Valency : 
3 and 5 ; Atomicity of Vapour : 2-4 ; Atomic Weight : 74-96 ; 
Melting Point, under pressure : about 817 ; Specific Gravity : 

History. In the form of its sulphides, realgar and 
wpiment, arsenic has long been known. The Greek alchemists 
employed it in their operations, and further investigations on 
it were carried out by the chemists of Islam in the early 


middle ages. They called the arsenic sulphides zarnlkh, and 
some of them regarded the element (which they knew how to 
extract from its sulphides) as a kind of mercury " Eastern 
Mercury." They supposed it to be a constituent of certain 
metals. It is interesting to note that the famous Muslim chem- 
ist ABU'L-QASIM AL-'!RAQI (thirteenth century A.D.) pointed out 
the close resemblance between zarnlkh ( As 2 S 3 ) and kuhl (Sb 2 S 3 ). 
ALBEBTUS MAGNUS (1193-1282) prepared metallic arsenic by 
heating orpiment with soap ; while arsenious oxide, As 2 O 8 , 
or " white arsenic," was made by roasting the sulphides and 
collecting the sublimate. * ' GEBBR ' ' (thirteenth century) gives 
the following instructions for the preparation of white arsenic 

" Arfnick [sulphide] is beaten to Powder and mujt then be 
boyled in Vinegar, and all its combujtible Fatnefs extracted, 
and it then dryed. Then R of Copper, calcined, lib. 1. Of 
Allom calcined | a pound, and of Common-Salt prepared as 
much as of the Allom. Mix the/e with your Arfnick prepared, 
and having ground all well together, moyjten the Mixture with 
dijtilled Vinegar (that it may be liquid) and boyl the Jame, as 
you did in Sulphur ; and then Jublime it in an Aludel (without 
an Alembeck) of the height of one Foot. Gather what a/cends 
white, denje, clear, and lucid, and keep it ; becauje it is 
fufficiently prepared for the Work." 

Occurrence. Small quantities of arsenic are found free in 
nature, but it occurs chiefly as realgar, As 2 S 2 , the red sulphide, 
orpiment, As 2 S 3 , the yellow sulphide, arsenical pyrites or 
mispickel, FeAsS, and cobalt glance, CoAsS. 

Preparation. When arsenical pyrites is heated out of 
contact with air the arsenic volatilizes and may be condensed 
and collected. Arsenic may be obtained from white arsenic, 
arsenious oxide, by heating with powdered charcoal. The 
mixture of white arsenic and charcoal is heated in a fireclay 
crucible, and the sublimate of arsenic collected in a conical 
iron receiver placed over the crucible 

As 2 8 + 3C = 2As + SCO. 

Properties. Arsenic resembles phosphorus in the fact 
that it exists in several allotropic forms. 


(i) Ordinary " metallic " arsenic or y-arsenic is a hard, 
brittle, greyish, crystalline substance which conducts heat and 
electricity well. It has a specific gravity of 5-73, is insoluble 
in carbon disulphide, and, in fact, in general physical properties 
is distinctly metallic. On heating it begins to volatilize at the 
temperature of boiling water ; at a higher temperature it 
rapidly changes it into a yellow vapour of density 150 (H = 1) 
at 860 and 75 at 1,800. These figures show that at 860 
the vapour consists of As 4 molecules (As = 75) and at 1,800 
of As 2 molecules. At intermediate temperatures some of 
each kind of molecule would be present. 

Metallic arsenic is used in the manufacture of shot, since 
an alloy of lead with about 0-5 per cent, arsenic is harder 
than pure lead and forms better shot. 

It is insoluble in water, but is oxidized by nitric acid to 
arsenic acid, H 3 AsO 4 . With sulphuric acid it yields arsenioua 
oxide. It takes fire spontaneously in chlorine, if finely 
powdered, forming arsenic trichloride, AsCl 8 . If heated in 
air or oxygen it burns with a characteristic bluish flame, 
forming arsenious oxide. 

(ii) " Amorphous " arsenic (^-arsenic) is obtained by vaporiz- 
ing " metallic " arsenic in a current of hydrogen and allowing 
the vapour to condense on the cold parts of the tube. It 
probably consists of very minute crystals of the ordinary 

(iii) Yellow arsenic (a-arsenic) (corresponding to yellow 
phosphorus) is made by cooling arsenic vapour in liquid air in 
absence of light. It is a yellow crystalline solid, soluble in 
carbon disulphide and readily oxidizing in the air. During the 
oxidation the arsenic glows, as does yellow phosphorus. It 
quickly changes into " metallic " arsenic on exposure to light. 


Arseniuretted hydrogen, arsenic hydride, or arsine* 
AsH 8 . Arsine is formed by the reduction of a solution of any 
soluble arsenic compound with nascent hydrogen. Under 
these conditions, however, it is mixed with excess of hydrogen, 
so that to obtain the pure gas other methods are used 


(i) Action of water on aluminium arsenide, on heating 
AlAs + 3H 2 O = A1(OH) 8 + AsH 8 . 

(ii) Action of dilute hydrochloric acid on zinc arsenide 

Zn 3 As 2 + 6HC1 = 3ZnCl 2 + 2AsH 8 . 

The gas may be collected by downward displacement. It is 
colourless and poisonous and has an offensive smell. It is prac- 
tically insoluble in water (cf. ammonia and phosphine). It 
burns with a bluish flame, forming water and arsenious oxide 
2AsH 3 + 30 2 = As 2 O 3 + 3H 2 O. 

When arsine is passed through a hot tube it is split up into 
arsenic and hydrogen, and the former element condenses as 
a black mirror on the cool parts of the tube farther on. (See 
tests for arsenic , p. 397.) 

Arsine is a reducing agent, and precipitates silver from a 
dilute solution of silver nitrate 

AsH 3 + 6AgNO 8 + 3H 2 O = 6Ag + 6HNO 8 + H 3 AsO 3 . 

Arsenious acid. 

Other hydrides of arsenic have been described, but they are 
ill- defined and unimportant. 

Arsenic trichloride, AsCl 8 , is a colourless fuming oily 
liquid (B.P. 130) formed by synthesis from its elements or 
by distilling a mixture of arsenious oxide, salt, and sulphuric 

As 2 O 8 + 6HC1 ^r 2AsCl 8 + 3H 2 O. 

Salt and sulphuric acid are used instead of concentrated 
hydrochloric acid since arsenic trichloride is partially decom- 
posed by water, first into a basic chloride and then into 
arsenious acid, hydrochloric acid being liberated 

(i) AsCl 3 + 2H 2 ^= As^-OH + 2HC1. 


(ii) AsA)H + H a O ^= H 3 AsO 8 + HC1. 

Compare the hydrolysis of phosphorus trichloride, which is 
practically complete, 

A better method of obtaining arsenic chloride was described 


by PARTINGTON (1929). He heated a mixture of arsenious 
oxide and sulphur chloride under a reflux cbndenser, and 
passed chlorine through it. Arsenic chloride was formed, and 
could be distilled off. 

If chlorine is passed into strongly cooled arsenic trichloride, 
it has been stated that the pentachloride is formed, AsCl 6 , 
though the observation has not been confirmed. 

Arsenious oxide, As 2 O 3 , "white arsenic" *is obtained by 
roasting the element itself, or its sulphides, in a current of 
air or oxygen. It occurs in three modifications 

(i) Amorphous or vitreous arsenious oxide, formed by care- 
fully and slowly condensing the vapour. In damp air it 
gradually passes into the crystalline octahedral form. It is 
the most soluble of the three forms. 

(ii) Octahedral arsenious oxide, the ordinary form, pro- 
duced when the vapour is condensed without special pre- 
caution. It is only slightly soluble in water. 

(iii) Rhombic arsenious oxide, formed by heating (i) or (ii) 
to about 200 for some hours. It is unstable. 

The aqueous solution of arsenious oxide has a slight acid 
reaction, owing to the formation of arsenious acid 

As 2 3 + 3H 2 ^z 2H 8 AsO 3 ^= 2H* + 2H 2 As0 3 x . 

This is a very weak acid and has never been isolated, although 
arsenites are known, e.g. Ag 3 As0 3 , silver arsenite. 

Arsenious oxide and the arsenites are good reducing agents, 
since they readily take up oxygen to form arsenic pentoxide 
and ar senates respectively. Sodium arsenite solution is often 
used for the estimation of iodine 

2Na 3 As0 3 + 2I 2 + 2H 2 = 2Na 8 AsO 4 + 4HI. 

Sodium arsenate. 

For this purpose, an excess of sodium bicarbonate is added 
to the sodium arsenite solution, to take up the hydriodic acid 
as formed. Alkali cannot be used, as it would react with the 

2NaOH + I 2 = Nal + NalO + H 2 0. 
Scheele's green is copper hydrogen arsenite, CuHAs0 8 . 
This and similar arsenical colouring matters were, and still are 


to some extent, used in wall-papers a dangerous habit, for 
if the walls are damp and the paper gets attacked by moulds, 
arsine (or an organic derivative of arsine) is set free. Cases 
of poisoning have often occurred from arsenical wall-papers. 

Arsenious oxide vapour at comparatively low temperatures 
(500-800) has a density corresponding to the formula As 4 O 6 , 
but as the temperature rises the vapour density gradually 
falls until it reaches a minimum value, corresponding to the 
formula As 2 O 3 , at about 1,800. 

Arsenious oxide and indeed all soluble arsenic compounds 
are extremely poisonous, 0-2 of a gram of white arsenic being 
a fatal dose. However, the system may become accustomed 
to this poison, which is often used to beautify the skin and 
improve the wind. 

Arsenic pentoxide, As 2 6 , is formed by acting upon 
arsenious oxide with concentrated nitric acid and igniting 
the arsenic acid so obtained 

2H 3 AsO 4 = 3H 2 O + As 2 O 6 . 

It is a white deliquescent crystalline solid, which dissolves 
in water to form arsenic acid. At a high temperature it splits 
up into oxygen and arsenious oxide. 

Arsenic acid, H 3 As0 4 , prepared as above, is a white 
crystalline solid containing water of crystallization. Its salts 
are the arsenates ; the only one of importance is disodium 
hydrogen arsenate, Na 2 HAs0 4 , which is used in calico-printing. 
Arsenic acid or arsenic pentoxide is sometimes used in organic 
chemistry as an oxidizing agent, e.g., in the preparation of 
quinoline. H 3 AsO 4 is or^oarsenic acid ; the corresponding 
pyro and meta acids are known, H 4 As 2 7 and HAs0 3 . 

Pyroarsenic acid is made by heating the ortho acid to 160 ; 
it dissolves in water with evolution of heat, reforming ortho- 
arsenic acid 

2H 3 As0 4 ^ H 4 As 2 7 + H 2 0. 

Meta-arsenic acid is formed when the ortho acid is heated 
to 200. It is a white substance that dissolves in water 
with evolution of heat, reforming orthoarsenic acid 
H 8 As0 4 ^= H 2 O + HAs0 8 . 


Magnesium ammonium arsenate, MgNH 4 AsO 4 , insoluble in 
water, yields the pyroarsenate on heating 

2MgNH 4 AsO 4 = Mg 2 As 2 O 7 + 2NH 3 + H 2 O. 

Arsenites and arsenates give a yellow precipitate with am- 
monium molybdate and excess of boiling nitric acid. (Cf. 
phosphates, p. 390.) 

Arsenic trisulphide, As 2 S 3 , or orpiment, is found 
naturally occurring. It was formerly used as a pigment, 
called auri pigmentum, whence the name " orpiment." It is 
obtained as a yellow precipitate by passing sulphuretted 
hydrogen through a solution of arsenious oxide in dilute 
hydrochloric acid 

2AsCl 3 + 3H 2 S ^= As 2 S 3 + 6HC1. 

It may be obtained in the colloidal state by passing sul- 
phuretted hydrogen through a hot solution of arsenious oxide 
in water. 

Arsenic pentasulphide, As 2 S 5 , is formed as a red 
precipitate by addition of hydrochloric acid to a solution 
of sodium thioarsenate 

2Na 3 AsS 4 + 6HC1 = GNaCl + 3H 2 S + As 2 S, 

Arsenic disulphide, As 2 S 2 , or realgar (the Arabic name 
of the substance), is a red mineral found in many parts of the 
world. It is used in making fireworks and in tanning. 

Thioarsenites and thioar senates. Arsenic trisulphide 
will dissolve in solutions of caustic alkalis, to form a mixture 
of an arsenite and thioarsenite 

As 2 S 2 + 4NaOH = Na 2 HAs0 3 + Na 2 HAsS 3 + H 2 0. 

Thioarsenites are also formed by dissolving arsenic tri- 
sulphide in a solution of an alkali sulphide, as in Group II of 
the analysis tables 

As 2 S 3 + 3(NH 4 ) 2 S = 2(NH 4 ) 3 AsS,. 

Addition of dilute acid to a solution of a thioarsenite pre- 
cipitates the arsenic as arsenic trisulphide. 

Solutions of thioarsenites will dissolve sulphur, giving 


thioar senates, e.g., Na 3 AsS 4 , which are also formed by dis- 
solving arsenic trisulphide in solutions of alkali polysulphides, 
such as yellow ammonium sulphides. The thioar senite is 
first formed and is then oxidized by excess of sulphur to the 
thioarsenate. The final product in Group II in analysis (see 
above) is therefore the thioarsenafe. 

Tests for Arsenic. Since the days of Paracelsus, arsenic 
compounds have been favourite tools of the poisoner, and as 
such a small dose is fatal the detection of arsenic becomes a 
matter of importance. Fortunately, there is no poison easier 
to detect, and the need for a systematic study of chemistry 
on the part of poisoners is becoming more and more evident ! 
The symptoms of arsenical poisoning have been prominent so 
many times in the newspapers of late years that there is no 
need for their description to be given here. 

The test usually employed for arsenic in criminal investi- 
gations is called MARSH'S test. All the apparatus and chemi- 
cals used must obviously be free from arsenic. Hydrogen is 
liberated from zinc and dilute sulphuric acid in a small flask, 
and a weighed quantity of the material to be tested is intro- 
duced. All the arsenic present, if any, is converted into 
arsine, and the mixture of arsine and hydrogen evolved is 
passed through a heated tube, where the arsine is decom- 
posed and arsenic deposited as a mirror on the cold parts of 
the tube beyond the flame. By comparing the mirror formed 
with those made under similar circumstances with known 
weights of arsenic, the weight of arsenic present in the material 
under investigation may be estimated. This method is sus- 
ceptible of very great accuracy. 

Arsenic in very small quantity is normally present in the 
human body. 

The arsenic mirror may be distinguished from a similar 
mirror of antimony (p. 400) by treatment with bleaching- 
powder solution, in which the arsenic is soluble and the 
antimony insoluble. 

For other tests for arsenic, see FENTON'S Notes on Qualitative 
Anal i /sis. 



Group in Periodic System : V ; Symbol : Sb ; Valency : 3, 
4, and 5 ; Atomicity of Vapour : 1-3 ; Atomic Weight : 121-8 ; 
Melting Point : 630-6 ; Specific Gravity : 6-71-6-86. 

History. Antimony and many of its compounds have 
been known from the most remote times. The naturally 
occurring sulphide was called by the Greek alchemists arijUjM 
(stimmi), stibi, larbason, or chalcedony. The modern names 
stibnite for antimony sulphide, Sb 2 S 3 , is derived from stibi t as 
is also the pseudo-Latin name for the element, stibium. 
DIOSCORIDES describes stibi as a shiny, brittle mineral, con- 
taining no earthy impurities ; if it is heated with red-hot 
charcoal it becomes like lead [that is, the sulphide is reduced 
to the metal]. 

Stibi was, and still is, used by the ladies of the East to 
darken the eyebrows. The Arabs called it kuhl ; thus the 
poet MUTANABBI (915-965) says : 

" How pale and dull the deepest black of kuhl 
Against the lustrous blackness of thine eyes ! " 

The name antimony is alleged to have arisen from the 
pseudonymous BASIL VALENTINE'S unfortunate experiment, 
about A.D. 1400, on the effect of the powdered metal as a tonic 
for monks ; the results were so disastrous that the substance 
was called anti-moine ! It is a pity that this story must be 
rejected as apocryphal, since the name wa?s in use at least 
300 years earlier. The chemist who wrote, about 1600, under 
the assumed name of Basil Valentine made, however, a 
thorough investigation of antimony and its compounds, and 
published his results in a wonderful book called the Triumphal 
Chariot of Antimony, which is, I think, the first monograph 
on a chemical element. All that was known of antimony 
until the end of the eighteenth century is clearly described 
in the Triumphal Chariot a masterpiece of chemical literature 
and obviously the work of a first-rate chemist. 

Occurrence. Small quantities of antimony occur natu- 
rally, but the element is generally found as stibnite, Sb 2 S 8 . 


Preparation. The crushed stibnite is heated with scrap 
iron in graphite crucibles 

Sb 2 S 3 + 3Fe = 2Sb + 3FeS, 

and the metal purified by fusion with potassium nitrate. It 
is also extracted by burning the sulphide to the oxide and 
reducing this with charcoal and sodium carbonate. 

(i) 2Sb 2 S 3 + 90 2 - 2Sb 2 3 + 6S0 2 . 
(ii) Sb 2 3 + 30 = 3CO + 2Sb. 

Properties. Antimony is a lustrous silvery metal, brittle 
and crystalline. Its specific gravity is 6-7 ; it melts at 630-6 
and boils at about 1,400. Antimony is stable in the air at 
ordinary temperatures, but will burn if heated, forming white 
fumes of antimony trioxide, Sb 2 O 3 , mixed with the tetroxide, 
Sb 2 4 . 

Dilute acids have no action on antimony, but it dissolves in 
the hot concentrated acids to give the chloride (with hydro- 
chloric acid), the sulphate (with sulphuric acid), or the 
hydrated oxide, " antimonic acid " (with nitric acid). 

Antimony is commercially important, since it will take a 
high polish and expands on solidification ; the latter property 
enables it to be cast in moulds. It is largely used in the form 
of alloys, such as Britannia metal (antimony 12 per cent., tin 
86 per cent., copper 2 per cent.), pewter (rather less antimony 
than in Britannia metal), and type-metal (lead 65 per cent., 
antimony 25 per cent., tin 10 per cent.). Type-metal and 
linotype metal are used for the type for printing. They are 
hard and can take very fine impressions, on casting. The 
best type (such as that made for ALDUS in Italy in the fifteenth 
century) is made from silver. 

Antimony, like phosphorus and arsenic, can be obtained in 
several allotropic forms. Yellow antimony is made by 
passing ozonized oxygen into liquid antimoniuretted hydrogen 
or stibine, SbH 3 , at 90. It resembles yellow phosphorus 
and yellow arsenic in being soluble in carbon disulphide, but 
differs from them in being, apparently, amorphous. 

Black antimony is obtained by the action of oxygen on 
liquid stibine at 40. It is amorphous. 


Explosive antimony was first obtained by GORE in 1858, 
by electrolysis of a strong solution of antimony trichloride 
in hydrochloric acid between a platinum cathode and anti- 
mony anode. A black shining deposit of antimony is formed 
on the cathode. It explodes violently when scratched with 
a hot glass rod. Gore showed that explosive antimony always 
contains about 6 per cent, of the trichloride together with 
traces of hydrochloric acid. It may be a solid solution of the 
trichloride in black antimony. 

Antimony hydride, antimoniuretted hydrogen, or 
stibine, SbH 3 , is prepared by the action of hydrochloric acid 
upon an alloy of magnesium and antimony containing 33 per 
cent, of the latter metal. The alloy may contain magnesium 
antimonide, Mg 3 Sb 2 , in which case the equation for the 
reaction would be 

M g3 Sb 2 + 6HC1 = 3MgCl 2 + 2SbH t . 

Mixed with hydrogen, it is obtained by the action of nascent 
hydrogen upon a solution of any soluble antimony salt. If 
the gas is passed through a hot tube, a deposit of antimony is 
obtained. (Cf. Marsh's test for arsenic, p. 397.) 

If pure stibine, obtained by the first method, is cooled by 
liquid air it forms a colourless liquid boiling at 17 and 
freezing to a white solid at 88. It is very soluble in certain 
organic solvents such as ether and petrol and especially 
carbon disulphide. 

With air or oxygen it reacts to give water and antimony and 
is a strong reducing agent. Thus it reduces nitric oxide to a 
mixture of nitrous oxide, nitrogen and ammonia, and gives a 
black precipitate of silver with silver nitrate solution. It has 
no basic properties, thus resembling arsine but differing from 
phosphine and ammonia. 

Antimony trichloride, SbCl 8 , is formed when antimony 
burns in chlorine or when antimony trisulphide is heated with 
concentrated hydrochloric acid 
(i) 2Sb + 3C1 2 = 2SbCl s . 
(ii) Sb 2 S, + 6HC1 ;= 2SbCl 8 + 3H 2 S. 

On evaporating the solution, crystals of the trichloride can 


be obtained. These are deliquescent and colourless and melt 
at 73. Antimony trichloride boils, without decomposition, 
at 223. It is decomposed by water, with formation of 
insoluble basic chlorides 

SbCl 3 + H 2 - SbOCl + 2HC1. 

SbOCl is antimony oxychloride or powder of Algaroth ; it is 
insoluble in water and is white in colour. The reaction 
SbCl 3 + H 2 O ^r SbOCl + 2HC1 is reversible ; thus anti- 
mony trichloride will dissolve completely in a little water, but 
on dilution the equilibrium is shifted very largely to the right 
and a white precipitate of the oxychloride is obtained. This, 
as would be expected, redissolves on addition of sufficient 
hydrochloric acid. A great excess of water carries the 
hydrolysis a stage further 

SbOCl + 2H 2 = Sb(OH) 8 + HC1, 

forming a hydratedantimonious oxide Sb(OH) 3 or Sb 2 s . 3H 2 0. 
Antimony pentachloride, SbCl 5 , is obtained as a colour- 
less or slightly yellow fuming liquid, B.P. 140, by passing a 
stream of dry chlorine through the trichloride. Although it 
may be boiled without decomposition, its vapour dissociates 
at higher temperatures into the trichloride and chlorine 

SbCl 5 ^z SbCl, + C1 2 . 

Water at ordinary temperatures decomposes it, yielding 
antimonic acid and hydrochloric acid. (Cf. phosphorus 
pentachloride, p. 382.) 

SbCl 6 + 4H 2 = H 3 SbO 4 + 5HC1. 

Compounds of antimony with bromine, iodine, and fluorine 
are also known. 

Antimony trioxide, Sb 2 3 , is made by hydrolysing the 
trichloride with hot water, and washing the precipitate with 
sodium carbonate solution until all hydrochloric acid is 

2SbCl 3 + 3H 2 O = Sb 2 3 + 6HCL 

It may also be prepared by strongly heating antimony in a 
current of steam. It is purified by sublimation. 


Antimony trioxide is a white powder which can be vaporized 
unchanged. Vapour- density determinations show that the 
vapour consists of Sb 4 O 6 molecules. It is almost insoluble in 
water, but dissolves in alkalis forming antimonites. It also 
dissolves in hydrochloric acid, yielding the chloride, but is 
insoluble in nitric and sulphuric acids. If boiled with a 
solution of potassium hydrogen tartrate it dissolves, forming 
potassium antimonyl tartrate or tartar emetic, often represented as 

2K(SbO)C 4 H 4 O fl .H 2 0. 

This compound is of great importance iri* medicine, and is 
also used in the dye industry as a mordant. 

Antimony tetroxide, Sb 2 4 . When the trioxide is 
heated in the air it burns, forming the tetroxide. This is a 
white solid which will dissolve in alkalis to form salts called 

Antimony pentoxide, Sb 2 5 , is made by acting upon 
antimony with strong nitric acid and carefully igniting the 
residue. It is a yellow powder, and on treatment with alkalis 
under suitable conditions can be made to yield antimoniates, 
e.g., sodium meta-antimoniate, NaSb0 3 . The acids corre- 
sponding to the pentoxide, viz., orthoantimonic acid, pyro- 
antimonic acid, and meta-antimonic acid, have all been prepared 
but are unimportant except that they form a further point of 
resemblance between antimony, arsenic and phosphorus. 

Antimony trisulphide, Sb 2 S 3 , occurs naturally as the 
grey mineral stibnite. It may be obtained as an orange red 
precipitate by passing sulphuretted hydrogen through a 
solution of antimony chloride acidified with hydrochloric 

2SbCl 3 + 3H 2 S ;= Sb 2 S 3 + 6HC1. 

The red form passes into the grey form if heated to 230 in 
an inert atmosphere. 

Thioantimonites and thioantimonates, corresponding to the 
thioarsenites and thioarsenates (p. 396), are known, and may 
be made in an analogous way. Ammonium thioantimonate, 
(NH 4 ) 3 SbS 4 , is made by dissolving antimony trisulphide ID 
yellow ammonium sulphide. 


Antimony pentasulphide, Sb 2 S 6 , is obtained as an orange 
precipitate by adding hydrochloric acid to a solution of a 

2(NH 4 ) 3 SbS 4 + 6HC1 = 6NH 4 C1 + 3H 2 S + Sb 2 S 6 . 

It is unstable, readily splitting up into the trisulphide and 


Group in Periodic System : V ; Symbol : Bi ; Valency : 2, 
3, and 5 ; Atomicity of Vapour : 1-2 ; Atomic Weight : 209-0 ; 
Melting Point : 269 ; Specific Gravity : 9-823. 

History. Bismuth was unknown to the ancient chemists 
or, if known, was considered to be a kind of lead. The name 
wismath was given by RULANDUS (Lexicon Alchemice Rulandi, 
Frankfort, 1612) to a certain metallic sulphide, but the name 
was afterwards applied to a new element, bismuth. " Wis- 
math " is derived from wis mat or wiesse Masse, white mass 
or metal ; the w was changed into b for the purpose of writing 
the name in Latin, from the alphabet of which language the 
letter w is missing. Bismuth and its compounds were studied 
by PARACELSUS and AGRICOLA and more especially by BASIL 
VALENTINE, but BERGMANN (eighteenth century) first care- 
fully distinguished between it and antimony, tin, and lead. 
The French chemist LEMERY, 1694, made a large fortune by 
selling " le seul Magistere de Bismut " (bismuth oxynitrate 
or pearl white) as a cosmetic to the ladies of the Court. 

Occurrence. Bismuth occurs chiefly as the native metal, 
but bismuth ochre (Bi 2 O 3 ) and bismuthite or bismuth glance 
(Bi 2 S 3 ) are also found in small quantities. 

Preparation. The native metal is melted and the molten 
bismuth allowed to run off from the earthy residue. If it is 
required to extract the metal from the ores, these are first 
roasted and then reduced by smelting with coke, scrap iron, 
and a flux. The metal is then purified by fusion with salt- 
petre, which oxidizes the impurities. A convenient method of 
obtaining the pure metal is to reduce the carefully purified 
oxide, e.g. with potassium cyanide. 


Properties. Bismuth looks very much like antimony ; 
it is a lustrous white metal with a red reflex. It is crystalline 
and brittle and is readily powdered. Its specific gravity ia 
9-8 ; it melts at 269, and boils at 1,450-1,500. It is a poor 
conductor of electricity (metals as a rule conduct well, while 
non-metals do not). 

Bismuth is stable in air at ordinary temperatures, but if 
heated strongly will burn with a bluish flame, forming the 
trioxide Bi 2 3 . It is dissolved by dilute nitric acid, yielding 
bismuth nitrate, Bi(NO 3 ) 3 . Other dilute acids have no action 
on it. Concentrated sulphuric acid acts upon it on heating, 
forming a basic sulphate or the normal sulphate according to 
the conditions. 

Bismuth is commercially important on account of the 
valuable alloys it forms. As a rule, these alloys are very hard, 
expand on solidification, and easily melt. The commonest is 
Wood's metal, which consists of bismuth 4, lead 2, cadmium 2, 
and tin 1 , and melts at 60-68. Others are Rose's metal (79), 
9 bismuth, 5 lead, 4 tin ; Newton's metal (94), 8 bismuth, 
5 lead, 3 tin ; Lipowitz's Alloy (70), 15 bismuth, 8 lead, 8 tin, 
3 cadmium. These easily fusible alloys are used for fuse 
wires, safety-valves for boilers, automatic fire-sprinklers, and 
so on. For the latter purpose it is advisable not to have the 
melting-point of the alloy too low. In the hot summer of 1921 
there were many reports of automatic sprinklers suddenly 
starting work unnecessarily ; the heat of the sun had been 
sufficient to melt the alloy which controlled them. The 
" magic spoons " that disappear when used to stir a cup of 
tea are made of a low melting-point bismuth alloy. 


Bismuth hydride, BiH 3 , has not yet been isolated. It is 
probably formed in small quantity by the action of concen- 
trated hydrochloric acid upon bismuth magnesium alloy. 

Bismuth trioxide, Bi 2 3 , is a yellow powder obtained by 
heating bismuth nitrate 

4Bi(N0 8 ) 8 = 2Bi a O 8 + 12N0 2 + 30,. 


It is used in the pottery trade for making glazes. 
It has no acidic properties, but is, on the contrary, basic, 
and dissolves in acids, forming bismuth salts 

Bi 2 3 + 6HC1 = 2BiCl 3 + 3H 2 0. 

Bismuth pentoxide, tetroxide, and dioxide, Bi 2 5 , 
Bi 2 O 4 , and Bi 2 O 2 , are also known. The pentoxide will 
dissolve in fused caustic alkalis to form bismuthates, e.g., 
NaBi0 3 , which are used as oxidizing agents. 

Bismuth trichloride, Bid 3 , is made by heating bismuth 
in a stream of chlorine. It is a white crystalline solid melting 
at 227 and boiling at 447. It is hydrolysed by water, 
forming bismuth oxychloride, a white insoluble powder 

BiCl 3 + H 2 ^= BiOCl + 2HC1. 

The oxychloride will dissolve in hydrochloric acid. 

Bismuth pentachloride is unknown. 

Bismuth nitrate, Bi(N0 3 ) 3 . In the form of its penta- 
hydrate, Bi(N0 3 ) 3 /5H 2 0, this salt is obtained by dissolving 
bismuth in dilute nitric acid and evaporating the solution to 
crystallization. It will dissolve in a small quantity of water 
or in dilute nitric acid, unchanged, but is hydrolysed to a 
white insoluble basic or subnitrate by excess of water 

Bi(N0 3 ) 8 + 2HoO ;== Bi(-OH + 2HN0 8 . 

Large excess of water hydrolyses the subnitrate further, to 
the hydroxide 

H 2 ^z Bi(OH) 8 + HNO,. 

" Bismuth subnitrate " is used in medicine, in cases of 
diarrhoea, etc., and is still employed, as in Lemery's time, as 
a cosmetic. It has, however, the disadvantage from the 
latter point of view of going yellow on exposure to air and 
thus giving the fair user a somewhat jaundiced appearance. 

Bismuth sulphide, Bi 2 S 3 , is insoluble in alkali sulphidei 
and does not yield thiobismuthites or thiobismuthates. 


Bismuth carbonate. The white precipitate obtained on 
adding ammonium carbonate to bismuth nitrate solution is a 
basic bismuth carbonate. 


1. Show that the elements nitrogen, phosphorus, arsenic, antimony, 

and bismuth are properly classified together. 

2. Describe the preparation of phosphine, phosphorus pentachloride 
and phosphorous oxide. 

3. How is phosphorus manufactured ? Give an account of the 
history of this element. 

4. Discuss the allotropy of phosphorus. 

6. Write a short account of the evolution of the match. 

6. What is the basicity of orthophosphoric acid ? Explain the 
behaviour towards indicators of aqueous solutions of orthophosphates. 

7. What do you know of the history of arsenic and antimony ? 

8. Compare and contrast the properties of arsenic with those of 
(a) phosphorus, (6) antimony. 

9. Show the stupidity of poisoners who use arsenic. 

10. How is stibine prepared ? What are its properties ? 

11. Give an account of the history and uses of bismuth. 

12. Is bismuth * metal ? Give your reasons. 


TYPICAL ELEMENTS : Oxygen, Sulphur. 

Sub-group A : Chromium, Molybdenum, Tungsten, 

Sub-group B (similar to typical elements) : Selenium, 


Oxygen and sulphur, both non-metals, have many resem- 
blances to one another, as will appear from the descriptions 
of these elements and their compounds. Close similarities 
between oxygen and chromium simply do not exist ; diligent 
search will enable us to find certain points in which sulphur and 
chromium are related thus, they both form acidic oxides 
MO 3 (sulphur trioxide, SO 8 , and chromium trioxide, Cr0 3 ), 
while the sulphates, M' 2 SO 4 , are isomorphous with the chro- 
mates, M' 2 Cr0 4 . It cannot be denied, however, that from 
the point of view of chemical relationship, the Periodic System 
has many failures. 

Oxygen, sulphur, selenium and tellurium are much more 
closely related to one another in properties than they are to 
the other elements of the group ; in fact, sulphur, selenium 
and tellurium form a triad of elements the members of which 
really do resemble one another very closely. The study of 
selenium and tellurium, however, lies beyond the scope of 
this book. 





Group in Periodic System : VI ; Symbol : ; * Valency : 2 ; 
Atomicity : 2 ; Atomic Weight : 16-00. 

History. Oxygen is said to have been known to the 
Chinese of the eighth century A.D. It is fashionable to 
attribute the discovery of any important substance or process 
to the Chinese, since it is difficult to disprove such a theory. 
Up to the present no one has suggested that the Chinese were 
acquainted with the existence of istopes. 1 Oxygen was dis- 
covered before 1773 by SCHBELE and in 1774 by PRIESTLEY. 
Scheele 's results were, however, not published until 1777, so 

FIG. 85. Scheele's Apparatus for the Preparation of Oxygen. 

that the credit of the discovery is usually given to Priestley. 
Scheele obtained the gas by heating a mixture of nitre (potas- 
sium nitrate) and oil of vitriol (sulphuric acid), and called it 
fire-air. He also made it by heating red oxide of mercury. 
It was from the latter substance that Priestley obtained 
oxygen. He heated the mercuric oxide by means of a burn- 
ing-glass, or double convex lens, which brought the sun's rays 
to a focus upon it. Priestley and Scheele both interpreted 
their results in the light of the current theory of combustion, 

1 The Greek alchemists were in the habit of talking about tne two 
lead*. What remarkable prescience 1 Seriously, though, the habit 
of reading modern theories into the ideas of antiquity is a vicious 
tendency that requires severe treatment. 


Nat. Portrait Gallery. 



the Theory of Phlogiston. In IdsPhysica Subterranea (1669), 
BECHEB said, " [Combustible] metals contain an % infiammable 
principle which by the action of fire goes off into the air ; a 
metal calx is left." According to the phlogiston theory, then, 
metals consist of two things a calx, different in different 
metals, and an inflammable principle common to all metals. 
This inflammable principle was called phlogiston by STAHL 
(1723), and combustion was considered to consist in the 
liberation of phlogiston from the burning body. A metallic 
calx could be reconverted into the metal by heating it with a 

substance rich in phlo- 
giston, such as carbon. 

. . . Sun's Rays ^ T i > T * 

" ' Cavendish s discovery of 

inflammable air proved of 
great assistance to the 
phlogiston theory, for it 
Burning Glass enabled the theory to ex- 

y' plain satisfactorily another 

Calx of Mercury set of facts. When a metal 

such as zinc is dissolved in 
dilute sulphuric acid, in- 
^| Mercury flammable air (considered 

by Cavendish to be nearly 
FIG. 86. Priestley 's Apparatus pure phlogiston) is evolved 

for the Preparation of Oxygen. i , . i f . ^ > ^ 

and a solution left which 

on evaporation yields crys- 
tals of white vitriol. Chemists argued that if zinc calx 
were taken and dissolved in dilute sulphuric acid, the solution 
of white vitriol should be obtained without evolution of in- 
flammable, air, since this has already been got rid of in con- 
verting the zinc into zinc calx. Experiment showed that this 
deduction was borne out in practice. 

An inconvenient observation for the theory of phlogiston 
was that when a given weight of metal is converted into its 
calx, an increase in weight occurs, whereas a decrease would be 
expected since phlogiston has been lost. Phlogistian chemist 
overcame this difficulty by assuming that phlogiston is lighter 
than air, and that therefore when it combines with a sub- 




stance tries to lift it as a balloon lifts a weight. Hence when 
a substance loses phlogiston it becomes heavier. % This theory 
received support from the extreme lightness of Cavendish's 
inflammable air. 

These facts are sufficient to show us that the phlogiston 
theory was a remarkably good one, correlating facts which 
were otherwise apparently disconnected, and predicting 
results that were afterwards verified by experiment. Let 
us imagine ourselves to be chemists of the eighteenth century, 
and attempt to explain Priestley's experiment on " red calx of 
mercury " in the light of the phlogiston theory. 

Calx of mercury we assume to be mercury minus phlogiston. 
We have heated this in a glass cylinder containing air, and 
have obtained mercury. This means that phlogiston must 
have been taken up by the calx ; where can this phlogiston 
have come from ? Obviously from the air in the cylinder. 
But the air in the cylinder is ordinary atmospheric air does 
this contain phlogiston ? Certainly, for fires all over the 
world are constantly liberating phlogiston and turning it into 
the air. Admitting this, what should be the properties of the 
residual air in the cylinder ? Clearly this air has lost phlo- 
giston, and is therefore " de-phlogisticated " ; it should there- 
fore be able to take up more phlogiston than the same volume 
of ordinary air in other words, (a) things should burn in it 
more brightly, since they can give up their phlogiston more 
readily, and (b) it should support life longer. Both of these 
results were actually obtained by Priestley, who therefore 
gave the name dephlogisticated air to his new gas, though the 
arguments he employed were not quite the same as those we 
ourselves have made above. 

However, the fact that an increase of weight occurs on 
calcination was felt by many chemists to be a serious objec- 
tion to the phlogiston theory, in spite of the ingenious expla- 
nation mentioned above. ANTOINE LAURENT LAVOISIER in 
1774 heated tin in a closed flask containing air, and found that 
the flask as a whole did not increase in weight. The tin> 
however, was converted into tin calx, and on opening the 
flask air rushed in. Lavoisier now weighed the flask and 


contents again, and made the important observation that the 
difference of weight between the tin and tin calx was equal to 
the weight of the air that rushed in when the flask was 
opened. Further experiments convinced Lavoisier that only 
a part of the air is concerned in calcination, and he therefore 
concluded that the air contains at least two different gases, 
only one of which is absorbed by the metal during calcination. 

In the autumn of 1774 Priestley dined in Paris with Lavoi- 
sier at the house of a mutual friend, Lord Shelburne, and 
described his discovery of dephlogisticated air. Lavoisier at 
once realized that Priestley's gas was probably that con- 
stituent of the air which combined with metals during com- 
bustion, and devised an experiment to test this idea. 

He confined mercury in a glass retort provided with a long 
neck communicating with air in a bell- jar placed in a trough 
of mercury. The level of the mercury in the bell- jar was 
noted by means of a strip of gummed paper. The mercury in 
the retort was heated for several days to a temperature just 
below its boiling-point. Lavoisier found that the mercury in 
the retort became covered with a red powder and that the 
level of the mercury in the bell- jar rose. After some time, no 
more red powder seemed to be forming and the mercury in the 
bell- jar stopped rising. At this stage Lavoisier noted the 
decrease in volume of the air in the bell- jar and found that 
it was approximately one-fifth. The residual air would not 
support combustion or life, so he called it azote. On heating 
the red calx or powder formed in the retort, he obtained a 
volume of gas equal to the diminution in volume of the air in the 
bell-jar ; this gas was of course Priestley's dephlogisticated 
air, and supported combustion extremely well. Lavoisier 
showed further that the increase in weight of the mercury in 
the retort during the above experiment was exactly equal to 
the loss in weight of the air in the bell- jar, and that a gas 
exactly like ordinary air could be produced artificially by 
mixing 1 volume of " dephlogisticated air " with 4 volumes 
of azote. 

His first conclusion was that air consists of two gases, azote 
and dephlogisticated air, in the proportion by volume of 4 to 1 j 


in combustion, only the, latter gas is concerned) and this combines 
with the burning substance, hence the increase in weight. The 
difference between the two theories of combustion is therefore 
as follows 

Phlogiston Theory : Metal = Calx + Phlogiston. 

Lavoisier's Theory : Metal + dephlogisticated Air = Calx. 

To mark the difference, Lavoisier re-named dephlogisticated 
air, calling it at first eminently breathable air, but afterwards 
oxygen (acid-producer) since he found that when non-metals 
were burnt in the gas and the products of combustion dis- 
solved in water the resulting solutions were acid. He con- 
cluded further that oxygen was an essential constituent of all 
acids, in which he was not quite right. 

Lavoisier's work forms the foundation of modern chemistry. 
On it Dalton was able to build his atomic theory, and from 
that time chemistry has never looked back. In 1869 WURTZ 
expressed the feeling of many chemists when he said, " La 
Chimie est une science francaise ; elle fut constitute par Lavoi- 
sier d' immortelle memoire." Lavoisier, however, was not 
appreciated by the Revolutionists, who in 1794 guillotined 
him, with the comment, " La Eepublique ria pas besoin de 

Occurrence. Oxygen forms 21 per cent, by volume and 
23 per cent, by weight of the air. In water it is combined 
with hydrogen in the proportion by weight of 88-9 per cent, 
oxygen to 11-1 per cent, hydrogen. About 50 per cent, of the 
earth's crust is composed of oxygen, and it is an essential 
constituent of all li ving matter. It is probably present in the 
sun and in other stars. 

Preparation. Oxygen may be made by heating mercuric 
oxide, manganese dioxide, nitre, barium peroxide, and many 
other substances. Potassium chlorate is usually employed. 
If this salt is heated to about 360 it melts and oxygen is 
evolved (Fig. 87) ; after a time the gas stops coming off and 
the liquid thickens, owing to the formation of potassium 
perchlorate, KC10 4 . 

8KC10, = 5KC10 4 + 3KC1 + 20,. 



If at this stage the temperature be raised to about 630, 
the potassium perchlorate is split up into potassium chloride 
and oxygen 

KC10 4 = KC1 + 20 2 . 

Omitting the intermediate stage, we can therefore represent 
the action of heat on potassium chlorate as 

2KC10 3 = 2KC1 + 30 2 . 

In the laboratory, it is usual to add a catalyst to increase 
the speed of the reaction. Manganese dioxide is generally 



FIG. 87. 

used, although many other substances will do equally well. 
The oxygen ia then evolved in a steady stream at a tempera- 
ture lower than the melting-point of the potassium chlorate 
(340 C.). 

Oxygen may be conveniently prepared in the laboratory 
by allowing water from a dropping -funnel to drop on to 
sodium peroxide in a flask fitted with a cork carrying a 
delivery-tube (Fig. 88) 

2Na 2 2 + 2H 2 = 4NaOH + 2 . 

The trade name for sodium peroxide is oxone. 

Pure oxygen is prepared by Baker's method electrolysis of 
a dilute solution of barium hydroxide, oxygen coming off from 
the anode. 

Commercially, oxygen is prepared by the fractional distilla- 



tion of liquid air. Since oxygen boils at 182 and nitrogen 
at 196, when liquid air is allowed to boil a partial separation 
can be effected. 

Oxygen is put on the market compressed in steel cylinders 
at a pressure of about 120 atmospheres. 

Properties. Oxygen is a colourless gas with no taste or 
smell. It is slightly heavier than air. It dissolves sparingly 

FIG. 88. 

in water, 1 litre of water dissolving about 50 c.c. of oxygen 
at C. It is, of course, the dissolved oxygen, and not the 
combined oxygen, in water that fish breathe. Oxygen readily 
supports combustion, the products being called oxides. 

Oxides are classified as follows 

(i) Acidic oxides, (ii) Basic oxides, (iii) Peroxides and 
Suboxides. (iv) Neutral Oxides. 

(i) Acidic Oxides. When non-metals are burnt in oxygen, 
the oxides obtained will generally unite with water to form 
acids, e.g., 

P 2 ft + H 2 = 2HP0 8 , meta-phosphoric acid. 
N t O fi 4- H a O = 2HNO 8 , nitric acid. 


S0 8 + H 2 = H 2 S0 4 , sulphuric acid. 
SO 2 + H 2 O = H 2 SO 3 , sulphurous acid. 

They are therefore called acidic oxides. Note that they are 
not acids they form acids when they unite with water. They 
may therefore be called acid anhydrides ; thus SO 8 is sulphur 
trioxide or sulphuric anhydride. 

Certain metallic oxides are acidic, e.g., Mn 2 7 , manganese 
heptoxide, and Cr0 8 , chromium trioxide. 

(ii) Metallic oxides are, however, usually basic oxides, and 
are true bases (contrast acidic oxides and acids), since they 
will react with acids to form salts + water. Some of them 
are capable of combining with water to form hydroxides, 
which are still bases 

CaO + H 2 O = Ca(OH) 2 , calcium hydroxide or slaked 

Note that hydroxides are not the same as hydrates. A 
hydroxide contains the hydroxyl group OH, while a hydrate 
is merely a loose compound of a substance with water, e.g., 
copper sulphate crystals, CuSO 4 .5H 2 O, are copper sulphate 
penta-hydrate ; BaCl 2 .2H 2 O is barium chloride dihydrate. 
In a solid hydrate, the water is generally present as water of 

Some basic oxides are soluble in water, giving alkaline 
solutions owing to the formation of the hydroxides, which 
yield OH 7 ions 

Na 2 O + H 2 O = 2NaOH ^r 2Na + 2OH 7 . 

The solution of a basic oxide, insoluble in water, by a dilute 
acid is easily explained in terms of the ionic theory : 

Suppose we have some zinc oxide in contact with water. 
We assume that, although the zinc oxide is apparently 
insoluble, it is in reality slightly soluble as the hydroxide, 
which ionizes into Zn" and 20IT 

ZnO + H 2 ^r Zn(OH), ^= Zn" + OH 7 + OH 7 . 

We should have then two equilibria, that between zinc 
hydroxide and zinc oxide and water on the one hand and that 
between un-ionized zinc hydroxide and zinc and hydroxyl 


ions on the other. When an acid, HX, is added, the new ions 
H* and X' are introduced, the former of which immediately 
combine with the OH' ions to form un-ionized water. In 
this way the equilibrium Zn" + 20H' ^ s Zn(OH) 2 is upset, 
and more zinc hydroxide ionizes, and this in turn causes more 
ZnO + H 2 to form zinc hydroxide. In this way the whole 
of the zinc oxide will be at last dissolved, if sufficient acid be 

Basic oxides may be prepared 

(a) By heating the metal in air or in oxygen. 

(b) By strongly heating the carbonate of the metal 

ZnC0 3 = ZnO + C0 2 . 

(c) By heating the hydroxide of the metal 

Cu(OH) 2 - CuO + H 2 0. 

(d) By heating the nitrate of the metal 

Pb(NO 3 ) 2 = PbO + 2N0 2 + oxygen. 

(e) By heating a peroxide of the metal, when the excess 

of oxygen is sometimes lost 

Pb 3 4 = 3PbO + oxygen. 

Not all these methods are applicable in every case ; some- 
times one method is more convenient and sometimes another, 
(iii) Peroxides and Sub-oxides contain respectively more 
and less oxygen than would be expected from the normal 
valency of the other element present. Peroxides are formed 
by both metals and non-metals. The commonest are 

Normal oxide. 

Sodium peroxide, Na 2 O 2 . . . . Na 2 O. 

Hydrogen peroxide, H 2 2 . . . H 2 0. 

Manganese dioxide, Mn0 2 . . . MnO. 

Lead peroxide, Pb0 2 .... PbO. 

Barium peroxide, Ba0 2 .... BaO. 

Nitrogen peroxide, N 2 O 4 . . . . N 2 O,. 

Chlorine peroxide, C10 2 . . . C1 2 0. 

Peroxides often yield oxygen on heating alone or when heated 
with strong sulphuric acid. With hydrochloric acid, they 
sometimes give hydrogen peroxide in the cold (e.g. Na,O a , 


Ba0 2 ), and on heating give oxygen (e.g. H 2 O 2 ) or chlorine 
(e.g. Pb0 2 + 4HC1 - PbCl 2 + 2H 2 + C1 2 ). A distinction 
has sometimes been drawn between peroxides and dioxides, 
but it is artificial and it is better to neglect it. 

Suboxides contain less oxygen than the normal oxide, e.g. 
carbon sub-oxide, C 3 2 . They are as a rule unstable. 

It often happens that an element has two or more normal 
valencies, in which case it may form two or more normal 
oxides, e.g FeO and Fe 2 3 , both normal basic oxides. In 
this case, the oxide containing a higher percentage of the 
element other than oxygen is called the -ous (from the Latin 
osus, richness) oxide, and the , other the -ic oxide ; thus FeO 
is f errors oxide and Fe 2 8 ferric oxide. Some peroxides 
behave as normal -ic oxides, e.g. when manganese dioxide, 
Mn0 2 , is treated with concentrated hydrochloric acid in the 
cold, it forms a salt + water 

Mn0 2 + 4HC1 =- MnCl 4 + 2H 2 O. 

It may therefore be called manganic oxide. On heating, 
however, it yields chlorine and manganous chloride 

Mn0 2 + 4HC1 = MnCl 2 + 2H 2 + C1 2 , 

and therefore is a peroxide. These facts serve to show that 
the classification of oxides given above is convenient only 
and not rigorous. 

(iv) Neutral Oxides. Certain oxides cannot be placed in 
any of the above classes they are neither acidic nor basic 
nor peroxides nor sub -oxides (in the usual sense that sub- 
oxides are unstable). Such are nitric oxide, NO, carbon mon- 
oxide, CO, and water, H 2 O. These are called neutral oxides. 

A fifth class of oxides might be made of those that show 
both acidic and basic properties. A common example is 
aluminium oxide, A1 2 8 . This will dissolve in acids to give 
a salt of aluminium 

A1 2 3 + 6HC1 = 2A1C1 3 + 3H 2 0, 
and also in caustic soda solution to give sodium aluminate 

A1 2 O 8 + GNaOH -= 2Na 3 A10 3 + 3H 2 O. 
It is therefore both basic and acidic. Such an oxide may be 


called amphoteric (not " anthropomorphic " as a Clifton boy 
recently wrote !). Other examples are zinc oxide, stannic 
oxide (Sn0 2 ), and lead peroxide (PbO 2 ). fn terms of the 
ionic theory, the phenomenon is explained in the following 
way. Aluminium oxide in the presence of water dissolves 
slightly, yielding aluminium hydroxide, A1(OH) 8 . This m 
solution ionizes in two ways 

(i) Al(OH) 3 ;=Ar- + 30H'. 
(ii) A1(OH) 3 ^ 3IT + A10 3 '". 

Both reactions occur simultaneously, the limiting factor being 
that the hydrogen and hydroxyl ions present are never in 
greater concentration than they are in pure water. It is 
clear that if an acid be added, the equilibria will be upset in 
such a way that the ionization (i) will be increased, while (ii) will 
be decreased ; whereas if a base be added, the reverse is true. 
Further Properties of Oxygen. Oxygen boils at 182 C. 
Liquid oxygen is a pale blue liquid, which is attracted by the 
magnet. Solid oxygen was first prepared in 1911 by Sir 
James Dewar by evaporating the liquid at extremely low 
pressures. Oxygen proved very difficult to solidify, and it 
was only by making use of the low pressures obtainable by 
the absorptive powers of charcoal cooled in liquid air that 
success was finally attained. The freezing-point of oxygen 
is 219 C. Oxygen is used commercially in welding by the 
oxy-acetylene flame, and medicinally in cases of difficult 


If the Chinese discovered oxygen, we may say that HOMER 
discovered ozone, although he is not usually considered the 
discoverer. It is, however, interesting to note that Homer 
mentions in the Iliad l the peculiar smell that occurs in the 
vicinity of a flash of lightning (or a " thunderbolt "), and to 
remember that in all the text-books VAN MABUM (1785) is 
called the discoverer of ozone because he noticed the same 
smell near an electrical machine in action, van Marum, 
ignoring the cat-skins on the machine, ascribed the smell to 
1 Bk. viii. 135. 



j 1 QQ QJ? ^- 

-* Oxygen. 

a new substance. Further researches by SCHONBBIN (1839) 
showed that this new substance was indeed the cause of the 
smell, and to it Schonbein gave the name of ozone, from the 
Greek $a), I smell. 

Occurrence. Ozone is said to occur in the air, and many 
thousands of people who smell the rotting seaweed on our 
shores remark on the health-giving properties of the ozone in 
sea-air. In point of fact, it is very difficult to prove that 
ozone is present in the lower atmosphere in any appreciable 
quantity: an argument 

seriously advanced was r / TRftJTnT N ' To Coil 

(a) the sky is blue ; 
(6) ozone is a blue gas, 
.'. (c) the air contains 
ozone. Ozone may be 
present in the atmo- 
sphere it probably 
is. Most things are. 

It is certainly true 
that ozone occurs in 
the air at a height of 
some 50 kilometres 
above sea-level, the 
amount changing with 
latitude and also with 
the weather conditions . 
The ozone shields the 
earth from intense ultra-violet radiation from the sun, 
which would otherwise cause serious effects. " It further 
causes the upper atmosphere at a height of 40-50 kilometres 
to be at a temperature in the neighbourhood of the normal 
boiling-point of water. This, in its turn, gives rise to the 
zones of abnormal audibility of sound from large explosions 
at a distance of some hundreds of miles from the explosion 
itself." (Nature, March, 1931.) 

Preparation. Ozone is usually prepared by passing the 
silent electric discharge through oxygen. It was for long sup- 
posed that the oxygen must be moist, and hence some chemists 

FIG. 89. Brodie's Ozone Apparatus. 



were of opinion that ozone was an oxide of hydrogen, but 
SHENSTONE (1907) in the laboratory at Clifton College showed 
that perfectly dry pure oxygen could be converted into 
ozone, which therefore must be an allotropic form of oxygen 
and not an oxide of hydrogen. This had long been suspected, 
of course. 

Various forms of ozonizer are used in the laboratory, one 
of the most convenient being that known as BERTHELOT'S, 
although it was invented by BRODIE in 1872. This is shown 
in Fig. 89. It consists of two coaxial tubes, in the inner of 
which sulphuric acid is placed. Oxygen may be passed 

Fia. 90. Siemens' Ozone Apparatus. 

through the annular space between the two tubes, and the 
outer tube is surrounded by sulphuric acid in a wide gas-jar. 

The two terminals of an induction coil are connected to 
two stout copper wires, one of which is placed in the inner 
tube of sulphuric acid and one in the sulphuric acid in the 
gas- jar. When the coil is working, a discharge passes from 
one Jot of sulphuric acid through the oxygen to the other, and 
some of the oxygen is converted into ozone. 

SIEMENS' ozonizer consists of two coaxial glass tubes, the 
outside of the outer one and the inside of the inner one being 
coated with tinfoil The oxygen is passed through the space 
between the tubes, and the two coats of tinfoil are connected 
one to one terminal of an induction coil and the other to the 



other terminal. The principle is thus exactly the same as 
that of Brodie's apparatus. 

The ozonair apparatus used in the Tube railways in London 
consists of a box containing plates of mica covered on each 
side with a coating of metal. Air is blown through the box 
and a discharge of electricity made to pass continuously 
between the plates. Another form of ozonizer used on a 
commercial scale is that invented by Siemens and Halske. 
This consists of an iron box containing water into which dip 
several porcelain tubes. Inside each of these tubes is another 
tube of aluminium. The iron box is connected to earth and 

Ozonised Air 





FIG. 91. Siemens-Halske Ozonizer. 

the aluminium tubes are now charged to a high potential. 
Air is passed over the aluminium tubes and through the 
porcelain tubes and is ozonized. 

Chemically, ozonized oxygen may be prepared by the action 
of nitric acid on ammonium persulphate, but all the methods 
hitherto described in this chapter convert at most 10-12 per 
cent, of the oxygen into ozone. Practically pure ozone (95 
per cent.) can be obtained by the electrolysis of acidulated 
water, using a high current density and a platinum anode 
the surface of which is scarcely as big as that of a pin's head. 
Under these conditions, ozone is evolved from the anode 

Properties. Ozone is a poisonous pale blue gas which 
can be condensed to an ultramarine liquid boiling at about 


112C. and freezing at 250 to a violet-black solid. Liquid 
ozone is fairly stable if pure, but readily explodes if organic 
matter is present. Ozone in the gaseous state*is more stable, 
but is catalytically decomposed by finely divided metals. 
Silver, however, is converted into silver oxide. (Cf . hydrogen 
peroxide on silver oxide.) Ozone is a strong oxidizing agent, 
and readily attacks rubber and other organic compounds. It 
bleaches indigo, and converts black lead sulphide into lead 
sulphate, which is white 

PbS + 40 3 = PbSO 4 + 40 2 . 

Ozone liberates iodine from a solution of potassium iodide, 
2KI + H 2 + O, = 2KOH + I a + 2 , 

but as many other oxidizing agents do the same, this cannot 
be used as a conclusive test for ozone. 

When ozone is passed over sulphur, a blue luminescence is 
shown, and sulphur dioxide and trioxide are obtained ; but 
the action is slow. 

One of the most characteristic properties of ozone is that 
it makes mercury " tail " or stick to glass, probably owing to 
the formation of a film of an oxide of mercury. The mercury 
can easily be cleaned afterwards by shaking with water. 

Ozone will not react directly with nitrogen, but moist 
ammonia is at once converted into ammonium nitrite and 
nitrate. This represents an oxidation of some of the am- 
monia to nitrous and nitric acids, which then combine with the 
excess of ammonia to form the corresponding ammonium salts. 
Ferrous salts are oxidized by ozone to ferric salts. Many 
organic compounds such as turpentine, ethylene, and benzene, 
combine directly with ozone to form ozonides, which are 
unstable compounds, often explosive 

H H H H 


C = C + 0, - > H C C H 


H H O O 


Ethylene. Ethylene ozonidat 

OZONE 425 

It will have been noticed from the above account of the 
properties of ozone that this gas can enter into reaction in 
three ways 

(i) Catalytic decomposition, 2O 3 >30 2 . 

(ii) In such a way that from one molecule of ozone one 
molecule of oxygen is liberated, e.g. 

2KI + H 2 O + O 3 -=*2KOH + I 2 + O t . 

(iii) In such a way that no oxygen is liberated, e.g. forma- 
tion of ozonides. 

Tests. (i) The smell, (ii) Liberation of iodine from 
potassium iodide solution (not characteristic), (iii) " Tail- 
ing " of mercury, (iv) Ozone will turn a piece of filter paper 
soaked in a solution of tetrabase (tetramethyl-p-p'-diamino- 
diphenylmethane) a purple colour ; hydrogen peroxide has 
no action on tetrabase, chlorine turns it blue, and oxides of 
nitrogen yellow, (v) Filter paper soaked in a solution of 
benzidine (another organic compound) is turned brown by 
ozone, red by chlorine and blue by oxides of nitrogen, while 
hydrogen peroxide has no action, (vi) Tincture of guaiacum 
is turned blue by ozone. 

Composition. Ozone consists of atoms of oxygen com- 
bined together in groups of three, that is, the molecules of 
ozone are triatomic (O 8 ) while those of oxygen are diatomic 
(0 2 ). Ozone is said to be an allotropic form of oxygen. 
When an element exists in more than one form, those forms 
being physically and often chemically distinct, it is said to 
exhibit allotropy, and the different forms are called allotropes 
or allotropic forms. Since in this case allotropy is shown to 
be accompanied by a different molecular structure, it is possible 
that allotropy is to be accounted for (a) by variations in the 
atomicity of the molecules of the allotropes, or (b) by varia- 
tions in the arrangement of the atoms in the molecules, 
whether the atomicity is the same or different. Allotropy is a 
common phenomenon. (See under sulphur and phosphorus.) 

The proof of the formula of ozone was carried out in a very 
elegant manner by SOBET. Soret enclosed two equal volumes 
of the same specimen of ozonized oxygen over water 




in two graduated tubes (Fig. 92). Into one tube he in- 
troduced a little turpentine, which absorbed the ozone 
without liberation of oxygen ; the decrease in volume there- 
fore is equal to the volume of ozone originally present. As 
an equal volume of the same specimen of ozonized oxygen 
was used in the second tube, this must contain the same 
volume of ozone as the first tube, and this volume is now 
known. The second tube was heated and then allowed to 
regain its original temperature, when it was found that an 

Increase in Volume, 



Decrease m Volume, 
Jg C-C.9 

After heating 
and cooling 

After addition ofa> 
little lurpenllnA. 

FIG. 92. 

increase in volume had occurred, and all the ozone was 
decomposed into oxygen. The volume of oxygen formed 
must be equal to the volume of ozone decomposed plus the 
increase in volume. 

Soret found that the decrease in volume in the first tube 
was exactly twice the increase in volume in the other. 
Hence, 1 volume of ozone gives 1J volumes of oxygen, .*. by 
Avogadro's Hypothesis, one molecule of ozone gives 1| 
molecules of oxygen. But the molecule of oxygen is O a 
/. that of ozone is 8 . 

The further question remains, what is the structural for 

OZONE 427 

mula for ozone ? If oxygen is uniformly bivalent, the 

structure of the ozone molecule must be /\ . From 


analogy with other substances, however, we might expect 
such a substance to be stable more stable indeed than 
oxygen itself but the reverse is true. Since one atom of 
oxygen is easily lost by the molecule of ozone, it is possible 
that this atom has a special position in the molecule. Such 
arrangements can be conceived of, if we assume that oxygen 
may act as a quadrivalent element in ozone. Many com- 
pounds are known in which oxygen is quadrivalent, so that 
there is no difficulty in assuming its quadri valence in ozone. 
We could then construct such formulae as 

O O 

= 0, /\ and /\ 

0=0 0=0 

Of these, the second would seem to be that most likely to 
lose an atom of oxygen, but at present there is no conclusive 
evidence on the point. 

Endothermic Nature of Ozone. When ozone is trans- 
formed into oxygen, a considerable amount of heat is evolved ; 
conversely, when ozone is formed from oxygen heat is ab- 
sorbed, therefore ozone is an endothermic substance in fact, 
when one gram-molecule of ozone is formed from oxygen, 
34,100 calories are absorbed. By the principle of Le Chatelier, 
therefore, we should expect ozone to be formed when oxygen 
is heated. However, if we pass a stream of ozonized oxygen 
through a hot tube, no ozone can be detected in the issuing 
gas. These facts seem to be contradictory to one another ; 
how are we to reconcile them ? 

In the first place, the reaction 2O 3 = 3O 2 is a reversible one, 
2O 3 s 30 2 . Hence at any given temperature and pressure 
there will be a certain definite equilibrium mixture of ozone 
atnd oxygen. At ordinary room temperature and atmospheric 
pressure, the percentage of ozone in the equilibrium mixture 
is quite inappreciable, and even at 1,300 C. it is only 0-15. 
But the ozonized oxygen from the ozonizer probably contains 


about 8 per cent, of ozone ; this is far too high a proportion 
for equilibrium, and the ozone will therefore decompose into 
oxygen until equilibrium is set up. This it does comparatively 
slowly, so that ozonized oxygen appears to be in equilibrium ; 
it is not really so. Now we know that a rise in temperature 
of 10 approximately doubles the rate of a chemical change. 
If therefore we heat the ozonized oxygen to say 420, the 
room temperature being 20 C., we shall increase the rate of 
decomposition of the ozone by 2 40 , and the ozone will reach 
its equilibrium concentration very quickly. But even at 
420, the percentage of ozone in the equilibrium mixture is 
too minute to be detected ; and hence the paradox. 

Uses. Ozone is an excellent germicide ; it is therefore 
used for sterilizing drinking-water (e.g. at Paris and Lille) and 
for the purification of air. The statement that it is beneficial 
to breathe ozone is untrue ; the gas has an irritating effect 
upon the mucous membrane. 

Oxidation and Reduction. When a substance combines 
with oxygen it is said to be oxidized ; reduction takes place 
when oxygen is removed from a substance. Thus 

2Cu + 2 = 2CuO 
represents an oxidation of copper to copper oxide, while 

PbO + C = Pb + CO 

represents the reduction of lead oxide to lead, carbon monoxide 
being evolved as well. 

Hydrogen is often used for effecting reduction, and the 
term was soon extended to cases in which the hydrogen 
merely added itself on to the substance concerned, and did 

/ H 

not actually remove oxygen ; thus, acetaldehyde, CH 3 . C/ , 

if treated with hydrogen under suitable conditions, ia con- 
verted into ethyl alcohol 

CH 3 . C*^ -f- H 2 = CH 3 .CH 2 .OH 

x o 

or C 2 H 4 + H, = C a H 6 0. 


Here the acetaldehyde is said to have been reduced ; the 
reverse process, namely, the conversion of ethyl, alcohol into 
acetaldehyde, is called oxidation, although no oxygen has 
been added to the alcohol hydrogen has been removed 

C 2 H 6 - H 2 = C 2 H 4 0. 

We may now define (i) oxidation, as the removal of hydro- 
gen from a substance, or addition of oxygen to a substance ; 
(ii) reduction, as the opposite of oxidation. The conversion 
of sulphuretted hydrogen into sulphur is therefore an 

H 2 S = H 2 + S. 

Hydrogen is a typical electropositive element and oxygen 
a typical electronegative element. The terms oxidation and 
reduction have therefore been still further extended in mean- 
ing, to include an increase in proportion of electronegative 
constituent in a substance (OXIDATION), and an increase in 
proportion of electropositive constituent (REDUCTION). 

(a) EeCl 2 + C1 2 = 2FeCl 8 (oxidation of ferrous chloride 

to ferric chloride). 

(b) K 2 Mn0 4 + chlorine = KC1 + KMn0 4 (oxidation of 

potassium manganate to potassium permanganate). 

(c) Cu + S = CuS (oxidation of copper to copper sul- 


(d) S + H 2 = H 2 S (reduction of sulphur to sulphuretted 


(e) MnCl 4 = MnCl 2 + Cl a (reduction of manganic chloride 

to manganous chloride). 

Oxidation and reduction may therefore take place although 
no oxygen or hydrogen is present. Note that every oxidation 
is necessarily accompanied by a corresponding reduction, and 
vice versa. Thus, in (c) above the sulphur is reduced by the 
copper, and in (d) the hydrogen is oxidized by the sulphur. 
In a great many cases, oxidation is accompanied by an 
increase in valency of the element or group oxidized, and 
reduction by a decrease. In (a) above, the valency of the iron 
is increased from 2 to 3 during the oxidation. In the case of 


electrolytes, it is true to say that oxidation involves increase 
of charge on the positive ion or decrease in charge on the 
negative ion, but to define oxidation in this way is- unsuitable, 
since non- electrolytes would not be included. 

Oxidizing and Reducing Agents. A substance that will 
bring about oxidation is an oxidizing agent, while one that 
causes reduction is a reducing agent. Common oxidizing 
agents are oxygen, ozone, hydrogen peroxide, chlorine, 
bromine, sulphur, nitric acid, chlorates, nitrates, peroxides, 
permanganates, dichromates. Common reducing agents are 
hydrogen, sulphuretted hydrogen, mixtures producing 
" nascent " hydrogen, hydrogen iodide, sulphur dioxide, 
carbon, zinc dust, stannous chloride, and many organic 


1. Write an essay on " chemistry without oxygen." 

2. What do you understand by oxidation and reduction T Mention 
the chief oxidizing agents and reducing agents. How would you test 
a given substance for (a) oxidizing powers, (6) reducing powers ? 

3. Write an account of the phlogiston theory of combustion, and 
estimate its value in the development of chemistry. 

4. Do you agree with the Public Prosecutor's statement that " La 
Republique n'a pas besoin de savants " ? 

5. What justification had Wurtz for saying that chemistry is a 
French science ? 

6. How does oxygon occur in nature ? 

7. Describe a laboratory method for the preparation of oxygen. 
What volume of oxygen, measured over water at 18 C. 740 mm., 
could be obtained from 30-0 grams of potassium chlorate ? 

8. Describe, with examples, the classification of oxides. 

9. How is ozone prepared (a) in the laboratory ; (6) commercially ? 
What are its commercial uses ? 

10. What is the formula for ozone, and how has it been ascertained T 

11. Ozone is an endo thermic substance ; how do you account for 
the fact that if ozonized air is passed through a hot tube, no ozone 
can be detected in the issuing gas ? 

12. A specimen of ozonized oxygen was found to diffuse 0*91 times 
as fast as pure oxygen under the same conditions of temperature and 
pressure. What is the percentage by volume of ozone in the specimen ? 

13. 1,000 c.c. of ozonized air were bubbled through potassium iodide 
solution. To react with the liberated iodine, 9-3 c.c. of N/50 sodium 
thiosulphate were required. Find the percentage by volume of ozone 
in the original gas. 


14. Mention ome characteristic tests for ozone. What is the action 
of ozone upon (a) mercury ; (6) lead sulphide ; (c) silver ; (d) ethylent* I 


History. Hydrogen peroxide was discovered by THENARD 
in 1818. He prepared it by the action of dilute acids upon 
barium peroxide, and showed that it contains twice the pro- 
portion of oxygen to hydrogen that is found in water. He 
therefore called it eau oxygenee. Its molecular weight was 
determined in 1893 by CARRARA, using the cryoscopic method ; 
he obtained the number 34, which corresponds to the formula 
H 2 2 . Pure hydrogen peroxide was first made by WOLFFEN- 
STEIN in 1894. 

Occurrence. Hydrogen peroxide occurs in minute quan- 
tities in the atmosphere. It is said that snow may contain as 
much as 0-0001 per cent. Traces of hydrogen peroxide have 
been found in plants. 

Preparation. Hydrogen peroxide is formed in small 
quantity (a) when ultraviolet light is passed through water, 
especially in the presence of oxygen ; (b) during electrolysis 
of acidulated water (probably as a result of the interaction of 
water and persulphuric acid) ; (c) by heating steam to an 
exceedingly high temperature ; (d) during the combustion of 
hydrogen and many organic compounds such as alcohol and 

It is frequently prepared from barium peroxide, Ba0 2 , or 
preferably barium peroxide dihydrate, BaO 2 .2H 2 0, by the 
action of an acid sulphuric or carbonic is generally used. The 
barium peroxide is made up into a thin paste with cold water, 
and this paste is then poured slowly into ice-cold water 
containing the calculated quantity of sulphuric acid to react 
completely with the barium peroxide, when 

Ba0 2 + H 2 S0 4 = BaS0 4 + H 2 2 . 

The barium sulphate is precipitated and may be filtered off. 

If carbonic acid is to be used, the paste of barium peroxide 

in ice-cold water is made rather thinner, and a stream of 

carbon dioxide is blown through. In this way the formation 


of an insoluble barium per carbonate, BaC0 4 , is avoided, 
whereas if the paste is added slowly to water through which 
carbon dioxide is blown (so that the latter is in excess) much 
of the percarbonate is precipitated and the yield of hydrogen 
peroxide is lowered. 

The aqueous solution of hydrogen peroxide prepared in this 
way may be evaporated on the water-bath until it reaches a 
concentration of about 50 per cent. It is then repeatedly 
fractionated under reduced pressure until a 99 per cent, 
solution of hydrogen peroxide is obtained. By cooling a 
little of this solution in a freezing-mixture of solid carbon 
dioxide and ether, crystals are obtained ; and if the bulk of 
the 99 per cent, solution is then cooled to 10 C. and 
inoculated with one of these crystals, a mass of transparent 
crystals of pure solid hydrogen peroxide separates, melting 
at -2 C. 

If the presence of salt is not deleterious, a solution of 
hydrogen peroxide may be conveniently prepared by adding 
sodium peroxide gradually to ice-cold dilute hydrochloric 

Na 2 2 + 2HC1 = 2NaCl + H 2 O 2 . 

The sodium peroxide may also be added to cold dilute 
sulphuric acid, most of the sodium sulphate crystallizing out. 
The residual liquid is fractionally distilled, the higher boiling- 
point fraction containing up to 30 per cent, of hydrogen 

JAUBERT'S method for ^obtaining a solution of hydrogen 
peroxide is to add citric acid to a solution of sodium perborate. 

Properties. Pure hydrogen peroxide at ordinary tem- 
peratures is a colourless syrupy liquid with an astringent 
taste and a strongly acid reaction to litmus. It is compara- 
tively stable, but explodes if brought into contact with 
platinum black or manganese dioxide. It sets fire to pow- 
dered carbon or magnesium and to cotton wool, but is said to 
have no action on reduced iron. Its specific gravity is 146 
and its boiling-point 85 at 68 mm. pressure. 

Hydrogen peroxide is an exothermic substance, but less so 


than water, therefore when it decomposes into water and 
oxygen heat is evolved 

2H 2 2 = 2H 2 + 2 + 44,800 calories. 

Aqueous solutions of hydrogen peroxide are usually kept in 
bottles lined with paraffin- wax, since they decompose in 
contact with glass. They are also readily decomposed, with 
evolution of oxygen, by finely divided or colloidal metals ; 
thus a colloidal solution of manganese containing 55 grams of 
manganese in 10,000,000 litres of water exercises a marked 
catalytic effect on the decomposition of hydrogen peroxide in 
aqueous solution. Colloidal platinum, copper, etc., have a 
similar action. 

2H 2 2 = 2H 2 + O 2 . 

Certain " organic catalysts " or enzymes act in the same 
way. The presence of traces of hydrocyanic acid, arsenic, 
etc., retards the action of the above catalysts, which are then 
said to be " poisoned." The decomposition of aqueous 
solutions of hydrogen peroxide is also retarded by the addition 
of alcohol, glycerol, calcium chloride, pyrogallol, or barbituric 
acid. These substances are therefore often added to com- 
mercial hydrogen peroxide solution to make it more stable ; 
they are good examples of negative catalysts. 

Hydrogen peroxide solutions are put on the market as 
" 10 volume," " 20 volume," and " 100 volume " ; these 
names refer to the volume of oxygen obtained by the de- 
composition of the hydrogen peroxide in 1 volume of the 
solution. The 30 per cent, solution is sold under the name 
of " perhydrol." 

Hydrogen peroxide is a powerful oxidizing agent ; it liberates 
iodine from potassium iodide solution, 

2KI + H 2 O 2 = 2KOH + I 2 , 
and converts lead sulphide into lead sulphate 
PbS + 4H 2 2 == PbS0 4 + 4H 2 0, 

just as ozone does. 

It will bleach many coloured substances, and as it is mild in 
action and leaves no injurious residue it is employed for 


bleaching silk, wool, and straw, as well as teeth and hair. It 
is unfortunate that the bright yellow colour given to hair by 
hydrogen peroxide is so easily recognizable. 

Owing to the readiness with which it loses oxygen it forms 
a good bactericide, and this property is made use of in hydro- 
gen peroxide mouth- washes and gargles. " Sanitas " and 
other disinfecting solutions also contain hydrogen peroxide. 

Action on Silver Oxide. Although hydrogen peroxide is an 
oxidizing agent, when it is added to silver oxide the latter is 
reduced to silver, and the hydrogen peroxide decomposed into 
water and oxygen 

Ag 2 + H 2 a - 2Ag + H 2 + 0,. 

This can be explained by the fact that when a molecule of 
hydrogen peroxide decomposes into water and oxygen, the heat 
evolved is greater than the heat of formation of silver oxide, 
which can therefore be split up into silver and oxygen. It is 
sometimes suggested that the " attraction " of the silver 
atoms for the one atom of oxygen and of the " water " in 
hydrogen peroxide for the other (H 2 O.O) is less than that of 
the two atoms of oxygen for one another to form a molecule 
of oxygen. This, however, is nothing more than an expres- 
sion of the fact of the reaction in other words ; it is not an 
explanation. Besides, if such attraction exists, why does not 
ozone reduce silver oxide to silver ? 

Potassium permanganate is readily decolourized by hydrogen 
peroxide in acid solution, both substances being reduced 
K 2 Mn 2 8 + 3H 2 S0 4 + 5H 2 2 

~ K 2 S0 4 + 2MnS0 4 + 8H 2 + 5O, 

This reaction may be used to estimate the strength of a 
hydrogen peroxide solution, but errors are likely to arise if the 
solution contains glycerol or other organic preservative, as 
many of these substances are themselves oxidized by potas- 
sium permanganate. 
Structure of the Hydrogen Peroxide Molecule. 

(i) On decomposition, 34 parts by weight of hydrogen 
peroxide give 18 parts by weight of water and 16 of 


.'. empirical formula = HO. 
(ii) Cryoscopic determination of molecular weight gives 34 

/. the formula = H 2 2 . 

(iii) The atoms in the molecule may be arranged in two 

H \ 

a. H H, or b. >O = O 

assuming in 6 that one oxygen is quadrivalent. 

In favour of the formula H H is the fact that ethy- 
lene combines with hydrogen peroxide to form ethylene glycol, 
which reaction is most easily expressed by the equation 

H H H 

\/ I 


II + I = I 


H H 

H H H 

Ethylene. Ethylene glycoL 

Similarly, when treated with diethyl sulphate (C 2 H 5 ) 2 S0 4 , 
hydrogen peroxide yields ethyl peroxide (C 2 H 5 ) 2 2 . Now if 

C 2 H 6 \ H\ 

this has the formula yO = (corresponding to 

C 2 Hg/ id./ 

for hydrogen peroxide) we should expect that on reduction it 
would give ethyl ether and water 

= + H 2 = >0 + H 2 

C 2 H 5 C 2 H 6 

Ethyl ether. 

Experiment shows that it actually gives ethyl alcohol, 

C 2 H 5 O 

which indicates that it has the formula, | 

C 2 H 5 
C 2 H 6 .0 

| + H, = 2C 2 H 6 .OH 

C,H 5 . Ethyl alcohol 



and therefore hydrogen peroxide would be | . 


However, the formula ^>0 = agrees better with the 


ready decomposition of hydrogen peroxide into water and 
oxygen, and also with certain physical properties. It is 
possible that ordinary hydrogen peroxide is an equilibrium 

H \ 
mixture of H H and >O = 


H \ 
H H ^= >0 = O. 


Similar cases of compounds reacting as though they had 
two different structures are common in organic chemistry, 
where the phenomenon is called tautomerism, and it has been 
shown that such compounds are generally equilibrium mix- 
tures of their two " tautomeric " forms ; hence it is reasonable 
to advance the same explanation in the case of hydrogen 
Tests for Hydrogen Peroxide. 

(i) The liberation of iodine from potassium iodide is not a 
conclusive test, since other oxidizing agents produce the same 
result ; it is, however, alleged that only hydrogen peroxide 
will liberate iodine from potassium iodide in the presence of 
ferrous sulphate. 

(ii) A dilute solution of potassium dichromate acidified 
with sulphuric acid gives a deep blue colour with hydrogen 
peroxide, owing to the formation of a perchromic acid (HCr0 5 
according to Moissan). This quickly decomposes with evo- 
lution of oxygen, but if ether is added it dissolves the per- 
chromic acid to a blue solution in which it is much more 

(iii) Tincture of guaiacum mixed with malt extract givus a 
blue colour with hydrogen peroxide. 

(iv) " Tetrabase " (p. 425) is not affected. 

WATER 437 

(v) Benzidine dissolved in alcohol is not affected. 

(vi) Titanium dioxide (Ti0 2 ) dissolves in sulphuric acid to 
a colourless solution. This is turned an orange colour by 
hydrogen peroxide, owing to the formation of TiO,, pertitanio 
anhydride or titanium trioxide. 


1. Describe the preparation and properties of pure anhydrous 
hydrogen peroxide. 

2. Mention three characteristic tests for hydrogen peroxide. How 
would you distinguish between a solution of this substance and (a) 
a solution of ozone ; (b) a solution of nitrous acid ; (c) a dilute solution 
of chlorine ? 

3. What is the formula for hydrogen peroxide, and how has it been 
arrived at ? What evidence have we as to the way in which the atorna 
are arranged in the molecule of hydrogen peroxide ? 

4. For what purposes is hydrogen peroxide used (a) in the labora- 
tory ; (6) commercially ? 

5. Hydrogen peroxide is an oxidizing agent. Mention facts in 
support of this statement, and explain the action of hydrogen peroxide 
on silver oxide. 

6. Calculate the strength of a solution of hydrogen peroxide which 
is exactly " 10-volume " at 15 C. 760 mm. 

7. 20 c.c. of a solution of hydrogen peroxide decolourized 18-5 c.o. 
of N/10 potassium permanganate. Calculate the strength of the 

8. 23-6 c.c. of a hydrogen peroxide solution were treated with 
excess of potassium iodide and dilute sulphuric acid. The resulting 
solution was made up to 250 c.c. 20 c.c. of this solution required 
15-4 c.c. of N/10 sodium thiosulphate. Find the strength of the 
original solution of hydrogen peroxide. 

9. How would you estimate the weights per litre of hydrogen peroxide 
and oxalic acid in a solution containing both these substances ? 


Occurrence. Pure water is never found in nature, but in 
a more or less impure state is widely distributed in enormous 


(i) Rain-water is the purest form of natural water. In 
country districts it contains only gaseous impurities, such as 


carbon dioxide, oxygen and nitrogen, which it has absorbed 
from the atmosphere, together with a little salt from the sea. 
Near towns, however, rain-water is much lees pure, and 
always contains soot and compounds of sulphur, such as 
sulphuric acid. 

(ii) River and spring water contains varying amounts of 
solid matter in solution, according to the ground in which it 
occurs. Salts of calcium render the water hard (p. 445). 

(iii) Mineral waters are waters containing unusual impuri- 
ties known or alleged to have certain curative properties. 
" Chalybeate " waters, for example, contain iron as ferrous 
bicarbonate, while the water of a certain spring at Epsom 
contains magnesium sulphate or " Epsom salt." Probably 
the most popular " mineral spring " in the world is the Well 
Zemzem at Mecca, the water of which is drunk by hundreds 
of thousands of pilgrims annually, and exported in bottles to 
all parts of the Muhammadan world. It is said to contain 
magnesium sulphate, arid, says SIR RICHARD BURTON, " It is 
apt to cause diarrhoea and boils, and I never saw a stranger 
drink it without a wry face." 

(iv) Sea-water contains about 3*6 per cent, by weight of 
solid matter in solution, chiefly the chlorides, sulphates and 
carbonates of sodium, potassium, calcium and magnesium. 
Salt forms about two- thirds of the total solid matter present. 

Composition. That water is a compound of hydrogen and 
oxygen was first shown by CAVENDISH in 1781. The volu- 
metric composition may be determined by explosion of a 
mixture of hydrogen and oxygen in a eudiometer tube over 
mercury (Fig. 93) ; it is found that 2 volumes of hydrogen 
require exactly 1 volume of oxygen, and, if the whole 
experiment is carried out at a temperature above the boiling- 
point of water, 2 volumes of steam are produced. 

Hence, by Avogadro's hypothesis, 

2 molecules of hydrogen + 1 molecule of oxygen give 
2 molecules of steam. 

.*. 1 molecule of hydrogen + molecule of oxygen give 
1 molecule of steam. 

Therefore the molecule of steam must be H,O. 



Very accurate experiments on the gravimetric composition 
of water were carried out by MOBLEY in 1895. A tube 
through which passed two platinum wires, for sparking, 
was fitted with two drying tubes containing phosphorus 
pentoxide, to prevent any loss of the water formed in the 
reaction. The whole was then evacuated and weighed. Pure 
dry hydrogen and pure dry oxygen were next admitted 

Fia. 93. 

(from weighed reservoirs) to the reaction tube through 
the drying-tubes, and the mixture ignited by a spark, the whole 
apparatus being kept cold by means of a freezing-mixture. 
Aft/er a suitable weight of the gases had been burnt the 
residual gas was pumped out and analysed and the tube and 
water weighed. The weights of the oxygen and hydrogen 
used were then known, and also the weight of water formed. 
Morley obtained as a mean of many results the ratio 



by weight of oxygen to hydrogen in water 7-9396 : 1. 

BUKT and EDGAR (1915) obtained the ratio 7*9387 : 1. 
The gravimetric composition of water is of additional 

importance since it gives at the same time the equivalent of 


Previous determinations of the composition of water by 
weight had been carried out by 
DUMAS and others, from 1820 on- 
wards. These determinations prac- 
tically all depended on passing 
hydrogen over heated copper 
oxide previously weighed, and 
measuring the weight of the 
water formed and the loss in 
weight of the copper oxide (Fig. 

Copper oxide + hydrogen = 
water + copper. 

The results so obtained were all 
more or less inaccurate owing to 
the experimental difficulties inherent 
in this method and to the fact that 
methods of obtaining absolutely 
pure hydrogen had not then been 

Morley's hydrogen was prepared 
by electrolysis, dried by phosphorus 
pentoxide, and absorbed by palla- 
dium (p. 183). On heating the 
"palladium hydride," pure hydrogen 
was evolved. The oxygen was 
made by the action of heat upon potassium chlorate. 

The molecular composition of liquid water is often assumed 
to be the same as that of steam, namely H 2 0, but as a 
matter of fact liquid water is a very complex substance 
and probably consists of an equilibrium mixture of H 2 0, 
"hydrol" (H 2 0) 2 , " dihydrol," and (H 2 O) 8 " trihydrol." 

FIQ. 94. Morley's 
Apparatus for the Compo- 
sition of Water. 



These may have the constitutions represented by the graphic 

H\ /H H\ xH 

0, >0 = 0< , and >0 0< . 
/ \ / \ 


H H 

According to ARMSTRONG, the last two of these " hydrols " 

H \ 
may exist each in more than one form, e.g. ^Q = 





, but the matter is highly speculative and 

cannot be discussed here. 

Ice is solid water, and melts at C. It exists in five or six 

FIG. 95. 

different forms of varying density. Ordinary ice is called 
Ice I and has a density of 0-92. 

Properties. Water is a practically colourless liquid, 
although it is said that in a layer of considerable thickness it 
has a bluish green colour. The difficulty of obtaining abso- 
lutely pure water in sufficient quantity makes any statement 
as to its colour open to question ; the fact that impure water 
is 'blue or green of course requires no proof, as it is obvious. 

At atmospheric pressure water freezes at C. and boils at 
100 C., these points of the thermometric scale being in fact 
fixed as the melting-point of ice and the boiling-point of water. 


The density of water is taken as unity, i.e. the weight of 
I c.c. of water at its point of maximum density (4 C.) is 
taken as the unit of weighty 1 gram. In the same way the 
capacity of water for heat is used in fixing the unit of heat ; 
thus the calorie is the amount of heat required to raise the 
temperature of 1 gram of water through 1 C. 

Water is an excellent solvent, and will dissolve practically 
every substance, at least in traces. It iar, therefore, very 
difficult to obtain pure. Pure water is usually prepared by 
distillation in apparatus made of block tin or of silver, with 
addition of a little potassium permanganate to oxidize any 
volatile organic matter. It may still contain carbon dioxide, 
from which it can be purified by redistillation, after addition 
of a little barium hydroxide. Water obtained in this way is 
called conductivity -water, as it is used in experiments on the 
electrical conductivity of aqueous solutions. 

Pure water is practically a non-conductor ; the purest 
specimen hitherto obtained had a specific conductivity 
(p. 132) of 0-4 X 10~ 7 reciprocal ohms. The concentration of 
hydrogen ions in water at ordinary temperatures is equal to 
that of the hydroxyl ions, and is about 10~ 7 gram-ions per 
litre, i.e. 10~ 7 grams of H" ions and 17 X 10~ 7 grams of 
hydroxyl ions, since the atomic weight of H = 1 and the 
molecular weight of the hydroxyl group is 17. 

Small as it is, this ionization of water is of great importance, 
since it leads to the hydrolysis of many salts of weak acids 
and bases in solution. Thus when sodium carbonate is 
dissolved in water the solution has an alkaline reaction, a 
phenomenon which can be explained by supposing that the 
water has partially bydrolysed the substance, with formation 
of caustic soda and carbonic acid ; the latter being a weak 
acid has scarcely any effect on litmus, whereas caustic soda 
U a strong base and turns litmus blue an alkaline reaction. 

Na 2 CO 3 + 2H 2 ^= 2NaOH + H 2 CO 3 . 

In terms of the ionic theory we have seen that the hydrolysis 
is really brought about by the ions of water, and will pro- 
ceed until so much carbonic acid has been formed that the 

WATER 443 

concentration of hydrogen ions from it is equal to the con- 
centration of hydrogen ions caused by the dissociation of 
water, H 2 - ^ H* + OH', that is, 10~ 7 gram-ions per litre. 
Many examples of the hydrolysis of salts will be found in 
this book ; hydrolysis, however, is not confined to salts but 
occurs with various other compounds as well. Thus phos- 
phorus pentachloride is hydrolysed in the presence of water, 
first to phosphorus oxychloride and then to phosphoric acid 

(i) PC1 6 + H 2 = POC1 3 + 2HC1. 
(ii) POC1 3 + 3H 2 = H 3 P0 4 + 3HC1. 

The hydrolysis of esters is important in organic chemistry. 
Esters form a class of compounds of which ethyl acetate, 
CH 3 . COOC 2 H 5 , is an example. When heated with water they 
are slowly hydrolysed, e.g. 

CH 3 .COOC 2 H 5 + H 2 ^=z CH 3 .COOH + C 2 H 5 OH. 

Acetic acid. Ethyl alcohol. 

The action is reversible ; the forward action may be made to 
go to completion by use of a large excess of water, or by 
adding caustic soda to take up the acetic acid as soon as it ia 
formed and hence prevent the reverse action. The hydrolysis 
of an ester (in absence of alkali) is greatly accelerated cata- 
lytically by the addition of a dilute acid. As any dilute acid 
will do, the catalysis is probably due to the hydrogen ions of 
the acid, and a method of determining the " strength " of an 
acid (i.e. its degree of ionization) is to measure its catalytic 
effect upon the rate of hydrolysis of an ester such as methyl 
or ethyl acetate. 

Hydration. Hydrolysis must be carefully distinguished 
from hydration. Many substances have the power of com- 
bining directly with water, to form compounds in which the 
water is more or less loosely held. These compounds are 
called hydrates and will generally give up their " water of 
hydration " on heating. In many cases hydrates are cry- 
stalline compounds which lose their crystalline form when the 
water is driven off : these are called crystalline hydrates and 
the water in them is called " water of crystallization." Copper 
sulphate crystals, for example, are copper sulphate penta- 


%drate,CuS0 4 .5H 2 0, while barium chloride crystals have the 
constitution BaCl 2 . 2H 2 . 

A characteristic property of crystalline hydrates is that at 
a definite temperature they show a perfectly definite vapour- 
pressure of water, which can easily be measured by introducing 
a little of the hydrate into the Torricellian vacuum. The 
vapour pressure of a hydrate increases on rise of temperature 
and falls with a lowering of temperature. 

Many crystalline hydrates effloresce on exposure to air ; that 
is, they lose some or all of their water of crystallization and 
fall to a powder. Experiment has shown that hydrates will 
effloresce if their vapour- pressure is greater than the partial 
pressure of water- vapour in the atmosphere. It sometimes 
happens, therefore, that a salt that is not efflorescent when 
the relative humidity of the air is high will effloresce when the 
air is drier. 

Deliquescence is a somewhat similar phenomenon but is not 
confined to hydrates, it may be shown by any substance 
which is very soluble in water. A deliquescent substance 
absorbs moisture from the air and gradually turns to an aque- 
ous liquid ; common examples are calcium chloride and 
caustic potash. Deliquescence occurs when the substance 
can form a saturated solution whose vapour-pressure is less 
than the vapour-pressure of water in the atmosphere, and will 
proceed until the solution is diluted to such an extent that its 
vapour-pressure is equal to that of the water in the air. The 
efficacy of a " drying-agent " for gases therefore depends 
largely upon its power of forming with water a solution (or a 
compound) having a low vapour -pressure of water, 

Thermal Dissociation. When steam is strongly heated 
it partially dissociates into hydrogen and oxygen 

2H 2 0;=2H 2 + 2 . 

This can be shown by blowing a rapid current of steam over 
a white hot platinum spiral and collecting the gases over water, 
when the steam condenses and a mixture of hydrogen and 
oxygen is left. At 2,500 about 10 per cent, of the steam is 

WATER 445 

Hardness. Water that will riot readily lather with soap 
is said to be hard. Soap consists mainly of sodium stearate, 
C 1T H 35 .COONa, the sodium salt of an organic acid, steario 
acid, C 17 H 35 .COOH. It is soluble in water, the solution 
having a " soapy " feel and lathering readily. If, however, 
soap is added to hard water, a curdy precipitate is obtained, 
and much more soap is needed to produce the lather. Analysis 
of the curdy precipitate shows that it consists chiefly of 
calcium stearate, although it may contain magnesium stearate 
as well. We may therefore conclude that hardness in water 
is due to the presence of soluble calcium or magnesium com- 
pounds which decompose the soap and thus prevent the 
formation of a lather until sufficient soap has been added to 
precipitate all the calcium and magnesium as stearates 

2C 17 H 35 COONa + Ca salt = (C 17 H 35 COO) 2 Ca + + 2Na salt. 

Calcium stearate, 
curdy precipitate. 

It is obvious that the hardness of water will depend on the 
weight of calcium or magnesium per litre and will be inde- 
pendent of the nature of the acid radical with which the metal 
is combined. The hardness of water is therefore usually 
expressed in grams of calcium per 100 litres (or in other units 

The chief salts which cause hardness are the bicarbonatee 
and sulphates. The former are present in water which has 
percolated through limestone or dolomitic (p. 244) rocks. 
Although calcium and magnesium carbonates are insoluble in 
pure water they dissolve in water containing carbon dioxide, 
owing to the formation of the bicarbonates, which are soluble 

CaC0 3 + H 2 O + C0 2 = Ca(HC0 8 ) 2 . 

Natural water always contains carbon dioxide, which it has 
acquired in its passage, as rain, through the air, or which it 
has absorbed from the soil, where carbon dioxide is con- 
stantly present. Hence if natural water flows over carbonate 
rocks it becomes hard. 

Hardness due to calcium or magnesium sulphate occurs 
only in those districts in which these compounds are present 


in the soil ; calcium sulphate is fairly widely distributed and 
hence is a common cause of hardness. 

Hardness may be estimated by titrating a measured volume 
of the water with a standard soap-solution until a " perma- 
nent " lather is obtained on shaking. (Permanency here 
means a duration of two minutes.) The soap-solution is 
standardized against a standard solution of calcium chloride. 

" Softening " of Hard Water. Water may be softened 
(that is, its hardness may be removed) by removal of the 
calcium and magnesium salts which it contains. In the case 
of hardness due to bicarbonates, merely boiling the water is 
sufficient, since the bicarbonates decompose on heating and 
the normal carbonates are precipitated 

Ca(HC0 3 ) 2 = CaC0 3 + CO 2 + H 2 0. 

Such hardness is therefore said to be temporary, as opposed 
to that caused by the sulphates. These cannot be removed 
by boiling the water and the hardness caused by them is 
called permanent. 

The softening of the water used in large towns is an impor- 
tant problem, since very hard water is in many ways objec- 
tionable not only for domestic use but also in industry. It is 
obviously impossible to soften the whole of a town's water 
supply by the process of boiling. Other methods are therefore 
employed. CLARK'S method for removing temporary hard- 
ness consists in adding to the water just sufficient lime to 
convert all the bicarbonate into normal carbonate (we may 
therefore call it the " homoeopathic " process !) 

Ca(HC0 3 ) 2 + CaO = 2CaC0 8 + + H 2 0. 

Care must be taken, of course, not to add too much lime, 
or the last state of that water will be worse than the first. 

Permanent hardness is removed by the addition of the 
calculated quantity of sodium carbonate to bring about the 

CaSO, + Na 2 CO, = Na 2 S0 4 + CaCO 8 + . 

A more modern method of softening hard water is the 
permutite process. " Permutite " consists of a sodium 

WATER 447 

aluminium silicate. When hard water is allowed to flow 
through a tube containing this substance the calcium and 
magnesium are retained as calcium and magnesium aluminium 
silicates ; the spent permutite can be regenerated by pouring 
on to it a strong solution of salt and washing out the calcium 
and magnesium chlorides so formed. 

FIG. 96. Formation of Stalactites and Stalagmites (Cox's Cave, 

Very soft water is unsuitable for a general water-supply* 
since not only does it taste " flat " (owing to absence of carbon 
Dioxide), but it also dissolves lead from the lead pipes through 
which it almost always flows at one part or another of its 
course. The action of the water on the lead in the presence 
of air results in the formation of lead hydroxide, Pb(OH) f , 


which is appreciably soluble. Lead salts are distinctly 
poisonous, and, as they do not pass out of the body, continued 
use of drinking-water containing lead would finally produce 
symptoms of lead-poisoning. With hard water, however, a 
coherent lining of lead carbonate is formed on the pipes and 
further solution of lead is therefore prevented. 

The "furring ' ' of kettles and the formation of " boiler-scale " 
are caused by the precipitation of calcium carbonate on 
boiling temporarily hard water. The formation of stalactites 
and stalagmites in caves in limestone districts (Cheddar, 
etc.) is due to the slow decomposition of calcium bicarbonate 
in the water which drips from the roof of the caves, with 
consequent deposition of calcium carbonate part being 
deposited on the roof (stalactite) and part on the floor when 
the drops fall (stalagmite). 

Recognition. Water may be recognized by its physical 
constants and also by the fact that it turns anhydrous copper 
sulphate, which is white, to the blue hydrated salt, CuS0 4 . 5H 2 0. 

Water -Analysis. The analysis of " water," i.e. detection 
and estimation of the impurities in it, is a subject in itself and 
cannot be adequately described within the limits of this book. 
Reference should be made to Button's Volumetric Analysis 
and to the technical books. 


The atmosphere. Air consists chiefly of a mixture of 
nitrogen and oxygen in the approximate proportion by 
volume of 4 to 1. That it is a mixture and not a compound 
is shown by the following considerations 

(i) The composition of the air, although fairly constant, 
does show slight variations, too large to be assigned to experi- 
mental error. 

(ii) Although the percentage composition by weight 
corresponds roughly to a formula N 4 (N 77 per cent. ; 
O == 33 per cent.), air cannot have this formula, for if it had 

4 X 14+16 

its vapour density would be = 36, whereas the 


experimental value is 14-4. 






(iii) This value, 14*4, is that which would be given by a 
mixture of 4 volumes of nitrogen, V.D. 14, with 1 volume 

of oxygen, V.Dr 16. 

(iv) The oxygen and 
nitrogen may be separated 
by mechanical means, e.g. 
fractional distillation of liquid 
air, diffusion, dissolving in 
water, etc. 

(v) The chemical and 
physical properties of air are 
those that would be expected 
of a mixture of nitrogen 
and oxygen in the propor- 
tions in which they occur in 

(vi) A gas closely resem- 
bling ordinary air may be 
made, without evolution or 
absorption of heat, by mixing 
4 volumes of nitrogen with 
1 volume of oxygen. 

The composition of dry air, 
free from carbon dioxide, 
may be found 

(i) Gravimetrically, by the 
method of DUMAS and 
BOUSSINGAULT. Air is sucked 
first through a bulb-tube con- 
taining caustic potash solu- 
tion to absorb carbon dioxide, 
then through a series of tubes 
containing pumice soaked in 
concentrated sulphuric acid, 
to remove moisture, and 
finally through a tube containing red-hot copper turn- 
ings ^weighed), into a weighed evacuated copper globe 
(Fig. 97). 

AIR 451 

The copper combines with the oxygen to form copper 
oxide, while the nitrogen passes on into the globe. 

Therefore, the increase in weight of the copper turnings 
is equal to the weight of oxygen which was mixed, in 
air, with the weight of nitrogen given by the increase in 
weight of the copper globe. 

It is hardly necessary to point out that the figure does 
not represent Dumas' actual apparatus (since Bunsen 
burners were not invented in 1843, the year of Dumas and 
Boussingault's experiment), but a modification suitable 
for the laboratory repetition of the experiment by an 
elementary student. 

(ii) V olumetrically , by enclosing the air over mercury in a 
graduated tube containing a copper wire which may be 
electrically heated to redness. The volume of air taken is 
noted, the wire heated to redness for a few minutes, and the 
apparatus allowed to cool. A decrease in volume will be 
observed, since the copper will have combined with the 
oxygen to form copper oxide, the volume of which is practi- 
cally the same as that of the copper wire itself and therefore 
may be neglected. After any necessary corrections for 
temperature and pressure, the volumetric composition of the 
air may therefore be calculated directly. The residual gas is 

Water -vapour is present in the atmosphere, and is an 
essential constituent as far as life is concerned. The per- 
centage of water- vapour in the air may be determined by a 
physical process (hygrometry , see textbooks of Heat), or 
chemically by aspirating a known volume of air through a 
weighed calcium chloride tube ; the increase in weight is due 
to absorption of the water-vapour of the air which has passed 
through. A cubic metre of air saturated with water- vapour 
at 25 C. contains 22-8 grams of water. The relative humidity 
of air at a given temperature is the ratio of the weight of 
water- vapour actually contained in 1 cubic metre of it to the 
maximum weight of water-vapour which 1 cubic metre of it 
could contain at that temperature. 

Carbon dioxide is present in the atmosphere in the proper- 


tion of 0-03 per cent, by volume. In spite of the fact that 
enormous quantities of this gas are continually being thrown 
into the air, from the combustion of carbonaceous fuels and 
the respiration of animals and plants, the percentage of 
carbon dioxide in the air remains remarkably constant. This 
is accounted for by the following facts 

(i) Carbon dioxide is soluble in water, so that any excess of 
the gas is dissolved by the sea. 

(ii) Carbon dioxide is the chief food of green plants. Plants 
do not, of course, breathe carbon dioxide ; they breathe 
oxygen, as do all other living things. They can, however, 
build up their tissues very largely from carbon dioxide and 
water ; this is an endothermic reaction and the necessary 
energy is supplied by the light of the sun. The active 
agent or catalyst in this process of carbon assimilation or 
photosynthesis is the green colouring-matter chlorophyll , 
a complicated organic magnesium compound related in 
structure to the red colouring-matter, hcemoglobin, of the 

The percentage of carbon dioxide in the air is most con- 
veniently determined by taking a definite volume of a standard 
solution of baryta-water, adding a few drops of phenol- 
phthalein and then forcing air through the liquid by means 
of a pump which delivers a known volume of air at each stroke. 
If the number of strokes necessary just to decolourize 
the solution is counted, the weight of carbon dioxide in a 
known volume of the air can be calculated. 

PETTENKOFER'S method is easy to carry out in the labora- 
tory as it does not require special apparatus. A Winchester 
bottle is filled with water and the water then poured into a 
measuring cylinder. This gives the capacity of the bottle. 
Fifty c.c. of a dilute standard solution of baryta are then put 
in the bottle, the cork replaced and the bottle well shaken. 
The carbon dioxide in the air in the bottle will react with some 
of the barium hydroxide 

(i) Ba(OH) 2 + CO, = BaC0 8 + H 2 0. 
If the residual liquid is now titrated, still in the bottle, with 


a dilute standard solution of oxalic acid, using phenolphthalein 
as indicator, the weight of unused barium hydroxide may be 

(ii) Ba(OH) 2 + H 2 C 2 O 4 = BaC 2 O 4 + 2H 2 0. 
Oxalic acid. 

The difference between the weight of barium hydroxide 
taken and the weight of that left is the weight of the barium 
hydroxide which has reacted with the carbon dioxide of the 
air in the bottle. From equation (i) above, it follows that 
171 grams of barium hydroxide react with 44 grams of carbon 
dioxide, hence the weight of the carbon dioxide in the air 
may be calculated. 

Other Gases, etc. Air contains traces of hydrogen, 
ammonia, nitric acid, oxides of nitrogen, hydrogen peroxide 
and ozone as well as bacteria, dust, spores and small quantities 
of various organic compounds. In addition, it contains about 
1 per cent, of the inactive gases, Helium, Neon, Argon, 
Krypton and Xenon, which are described in Chapter XVIII, 
p. 187. 

The average composition of the air in the latitude of London 
is as follows : 

Per cent, 
by volume. 

Nitrogen 77-32 

Oxygen 20-80 

Argon, etc 0-93 

Carbon dioxide ..... 0-03 

Water- vapour . . . . . .0-92 


Group in Periodic System : VI ; Symbol : S ; Atomic 
Weight : 32-06 ; Valency : 2, 4, or 6 ; Specific Gravity : 
1-98-2-06; Boiling Point: 444-6. 

History. The word sulphur shows us that this element 
and one of its most characteristic properties have been known 
for some thousands of years, for " sulphur " is the Latin form 
of the Sanskrit word sulvari or " enemy of copper." The 
Hindoos of 3000-2000 B.C. were therefore acquainted with 


sulphur and with the fact that it " destroys " copper when 
the two are heated together. To the Greeks and Romany 
sulphur was well known, since it occurs naturally in Sicily and 
other volcanic regions of the Mediterranean. The alchemists 
of Islam called it kibrit, and to them is due the theory that 
all metals are composed of kibrit and zlbaq (sulphur and 
mercury) a theory that was later modified into the 
phlogiston theory of Becher and Stahl. 

The properties of sulphur were thoroughly investigated in 
the Middle Ages, and many sulphur compounds described. 
The elementary nature of sulphur was shown by LAVOISIBK 
in 1777. 

Occurrence. Sulphur in the elementary state occurs in 
Sicily and Italy and, at a depth of about 700 feet, in Louisiana 
(U.S.A.). Combined as sulphates and sulphides, sulphur i 
very widely distributed. The chief sulphides are iron pyrites, 
JFeS 2 , copper pyrites, CuFeS 2 or Cu 2 S.Fe 2 S 3 , galena, PbS, 
zinc blende, ZnS, realgar, As 2 S 2 , orpiment, As 2 S 3 , and cinnabar, 
HgS. All these substances are of commercial importance, 
mostly on account of the metal they contain, although iron 
pyrites is used as a source of sulphur in the manufacture of 
sulphuric acid. The chief sulphate naturally occurring is 
gypsum, CaSO 4 .2H 2 0. 

Sulphur is a common constituent of living matter and i 
probably essential. 

Extraction. The native sulphur of Sicily is mixed with 
silica, limestone and other impurities. It is interesting, if 
somewhat disappointing, to find that the method used in its 
extraction is the same now as it has been for the last 1,000 
years ; only in detail have chemists been able to suggest any 
improvement. The sulphur ore is piled in heaps, called 
calcaroni, with the largest lumps at the bottom. Air-passages 
are left, and the floor slopes downwards so that the molten 
sulphur produced may flow out and be collected. The heaps 
of ore are covered with the residue from a previous operation, 
and are then set alight by the introduction from the top of 
burning wood. Part of the sulphur burns and melts the 
rest, which sinks to the bottom of the heap and flows out. 



It is collected in wet wooden boxes (Fig. 98). In this pro- 
cess, about 35 per cent, of the sulphur is lost by combustion. 

An improved kiln was invented by ROBERT GILL (1880) ; 
here the ore is heated in brick-work chambers arranged 
in groups of two or, more recently, six, in such a way that 
there is a minimum loss of heat and of sulphur. Even in 
the Gill kiln, however, there is a loss of 20-25 per cent. 

In 1891, a method for extracting sulphur by the use of 
superheated steam under pressure was patented. This had 
previously been suggested by PA YEN and GILL in 1867, but 
the practical difficulties 
at that time paoved 
insurmountable. Owing 
to the high price of coal, 
this method is not much 
used in Sicily, but is 
employed in the 

Refinement. The 
crude Italian and 
Sicilian sulphur is re- 
fined partly in the 
Romagna, but chiefly at 
Marseilles and Antwerp. p IG gg. 

The plant employed was 

invented by MICHEL in 1808 and improved by LAMY (1844) and 
DuJARDiN(1890). The crude sulphur is melted in iron pots 
(A, Fig. 99) and then flows into shallow iron retorts (B), 
where it is boiled. The vapour is passed into a large 
brickwork chamber (C), where it condenses (i) to a powder 
(flowers of sulphur), if the temperature of the walls of the 
chamber is below 100, or (ii) to a liquid when the tempera- 
ture rises above 114. 

The flowers of sulphur first formed are scraped out and sold 
as such. The liquid sulphur that later collects on the floor 
of the chamber is run out and allowed to solidify in cylindrical 
moulds, forming the so-called " roll-sulphur." 

Louisiana Process. Until the end of the nineteenth century 



practically all the sulphur on the market came from Sicily, 
but the discovery of a way to extract the sulphur in the 
Louisiana deposits resulted in a fall of some 45 per cent, in 
the sales of the Sicilian sulphur in 1912 and a still further 
fall in subsequent years. FRASCH (1903) made a boring of 
about 1 foot diameter through the overlying rock until he 
reached the sulphur at a depth of some 600-800 feet. Super- 
heated water was then blown down and the molten sulphur 

FIG. 99. Sulphur Refining. 

thus produced forced up by means of compressed air. A 
single well may produce as much as 500 tons daily. The 
sulphur obtained in this way is very pure (over 99 per cent.), 
and as the deposits are estimated at 40,000,000 tons, the supply 
of sulphur for the next century or so seems well assured. 
Large deposits have recently been discovered in Alaska, but 
they are not yet worked. 

Recovery of Sulphur from various Manufacturing Processes. 

(a) Manufacture of Coal-gas. Crude coal-gas contains 

sulphuretted hydrogen (from the sulphurous compounds in 


coal). It is passed over wet iron rust, when the sulphuretted 
hydrogen is removed 

2Fe(OH) 3 + 3H 2 S = Fe 2 S 3 + 6H 2 0. 

The iron sulphide so obtained is exposed to moist air for 
some time, when 

2Fe 2 S 3 + 6H 2 O + 30 2 = 4Fe(OH) 3 + 6S. 

The sulphur produced may be extracted with carbon 
disulphide, in which it is soluble, or may be burnt and thus 
converted into sulphur dioxide, which is commercially 

(b) From the residues of the Leblanc soda industry (CHANGE- 
GLAUS METHOD). See Manufacture of Sodium Carbonate, 
p. 204. [N.B. The Leblanc process is now obsolete.] 

The alkali-waste, containing calcium sulphide, CaS, in 
suspension in water, was treated with a stream of " chimney- 
gas " or " limekiln-gas " (carbon dioxide and nitrogen), 
when calcium carbonate and sulphuretted hydrogen were 

CaS + C0 2 + H 2 = CaC0 8 + H 2 S. 

The gas was collected in large gasometers and was then 
mixed with insufficient air for complete combustion and 
passed over heated iron oxide in a brickwork chamber called 
the Glaus kiln 

2H 2 S + 2 = 2H 2 O + 28. 

Sulphur was left in the kiln, and the iron oxide was un- 
changed ; it probably acted as a catalyst. Sometimes the 
sulphur was burnt to sulphur dioxide for use in the lead- 
chamber process for the manufacture of sulphuric acid. 

In this way nearly 100,000 tons of sulphur were recovered 
annually, and not only was much money saved but the moun- 
tainous heaps of alkali-waste no longer poisoned the air 
for miles around. 

Properties. Sulphur exists in a large number of allotropio 
forms. Ordinary sulphur is called a-sulphur or rhombic 
sulphur (since its crystals belong to the rhombic system). 

a- Sulphur. The naturally occurring form ; rhombic crystals 



melting at 114-5 C. Soluble in carbon disulphide, from which 
it crystallizes out in the same form. Specific gravity, 2-06. 
It is the stable form of sulphur at ordinary temperatures ; 
all other forms of sulphur change into it more or less 
quickly. If, however, the temperature is raised above 95-6 C. , 
a-sulphur becomes less stable than ^-sulphur, into which it 
therefore passes. 

ft '-Sulphur, monoclinic, or prismatic sulphur can be con- 
veniently prepared by melting a-sulphur in a crucible, allowing 
the liquid to cool until a thin crust has just formed, piercing 
this with two holes and pouring out the remaining liquid 

through one of these while air 
enters through the other. If the 
crust is now removed the crucible 
will be found to be lined with long, 
transparent, needle-like crystals of 
j8- sulphur. This form has a specific 
gravity 1-96 and melts at 119. 
Above 95-6 /^-sulphur is stable 
under atmospheric pressure, but 
below 95-6 it slowly passes into 
a-sulphur. a-sulphur, on the other 

handj is stable below 95 ' 6 ' but 
i ? * w unstable above. At 95-6, a- and 

a. Monoclinic. b. Rhombic. * 

p-sulphur are equally stable, and 

this temperature is called the transition-point for the two 

Here, then, we have an example of a substance that exists 
in two allotropic forms, either of which can be converted into 
the other by merely altering the temperature. These sub- 
stances are called enantiotropic (" moving in opposite ways "), 
as opposed to monotropic substances, such as phosphorus and 
iodine chloride. 

Phosphorus, for example, is capable of existing as yellow 
phosphorus and as red phosphorus. Now at all temperatures 
below the melting-point of the yellow form the latter is un- 
stable with respect to the red, into which it passes slowly at 
ordinary temperatures but more quickly on heating. Hence 



there is no transition-point between red and yellow phos- 
phorus, that is, no temperature at which both are equally 
stable, and below which one is the stable form while above it 
the other form is the stable one. A mere change of tempera- 
ture is sufficient, therefore, to change yellow phosphorus into 
red, but not the red into yellow : phosphorus is therefor 
called a monotropic substance. 

Equilibria between various Forms of Sulphur in Terms of the 
Phase Rule. The Phase 
Rule (p. 146) is P + F 
= C + 2. The constit- 
uent of all phases here 
is sulphur only ; hence 
= 1 and the rule 
becomes P + F = 3. 
If therefore we have 
three phases the system 
is non- variant, with two 
phases the system is uni- 
variant, and with one 
phase bivariant. Let us 
consider the phases a- 
sulphur, /S-sulphur, liquid 
sulphur, and sulphur 


96 115 1W T 

Fia. 101. Phase Rule Diagram for 

vapour. Below 95-6 we 

have a-sulphur and sulphur vapour (although the vapour 
pressure of a-sulphur is very small), that is, we have two 
phases and the system will therefore be univariant. In 
other words, if we fix the temperature, the pressure of sul- 
phur vapour will adjust itself to a definite value for that 

In figure 101, AB represents the vapour-pressure curve for 
a-sulphur, and as a-sulphur if heated quickly may be taken 
beyond the transition- point (95-6) and melted at 114-5, BO 
is the continuation of the vapour-pressure curve of this form 
as far as its melting-point. Above 95-6, however, a-sulphur is 
unstable, and changes slowly into ft- sulphur . If this change has 
taken place, then BE represents the vapour- pressure curve of 


^-sulphur from the transition-point up to the melting-point of 
this form, 120. At 95-6 we have the three phases a- sulphur, 
/J-sulphur, and sulphur vapour ; the system is therefore non- 
variant. EF is the vapour- pressure curve of liquid sulphur, 
and EG the curve showing the effect of pressure upon the 
melting-point of /3-sulphur. BG shows the effect of pressure 
upon the transition- point of a- and /?- sulphur, and CG the 
effect of pressure upon the melting-point of a-sulphur. 

We see therefore that /?- sulphur can exist in a stable form 
only within the limits shown by the triangle BGE. At G, 
a- and /^-sulphur and liquid sulphur are in equilibrium ; at E, 
/?-sulphur, liquid sulphur, and sulphur vapour ; at C, a-sulphur, 
liquid sulphur, and sulphur vapour, although at this point 
<z-sulphur is unstable with respect to ^-sulphur. B, C, E, and 
G all therefore represent non- variant systems. 

Other Allotropic Forms. Nacreous sulphur may be 
obtained by heating sulphur to 150, cooling the liquid to 
98, and scratching the sides of the containing vessel with a 
glass rod. It is a modification of /J-sulphur. 

Colloidal sulphur has been obtained from the sulphur which 
is precipitated by addition of sodium thiosulphate to cold 
concentrated sulphuric acid. 

Amorphous sulphur is present in small quantity in ordinary 
flowers of sulphur. It is insoluble in carbon disulphide and 
hence may be obtained by extraction of flowers of sulphur 
with this liquid ; the rhombic sulphur dissolves and amor- 
phous sulphur is left. It is a white powder. 

Plastic sulphur can be obtained by pouring molten sulphur 
into cold water. It is an elastic substance readily changing 
into the rhombic and amorphous forms. It is considered to 
be a supercooled liquid, that is, a liquid taken so quickly to 
far below its freezing point that it has not had time to crystal- 
lize. Glass is another example of a supercooled liquid. 

<f>-Sulphur is an orange-yellow crystalline form prepared by 
adding ice-cold concentrated hydrochloric acid to a cold 
solution of sodium thiosulphate and shaking the liquid 
with the organic liquid toluene (C 6 H 5 .CH 3 ). 

Milk of Sulphur is a white precipitate of sulphur obtained 


by adding dilute hydrochloric acid to a solution of calcium 
polysulphide (CaSj prepared by dissolving sulphur in boiling 
milk of lime. 

A-Sulphur and [t-Sulphur. When sulphur is melted, SA 
and Sju, are formed ; SA is of a pale amber colour and is mobile ; 
it is stable up to 160. Above 160 liquid sulphur contains 
chiefly S// which is dark in colour and viscous. It is said that 
Sft is formed only if the sulphur used is slightly impure. 
Other allotropic forms of sulphur have been described, e.g. 
Sp (" Engel's sulphur "), obtained by adding hydrochloric 
acid to sodium thiosulphate solution, filtering, extracting the 
filtrate with chloroform, and evaporating the chloroform 
extract. Sp is probably S. 

Sulphur boils at 444-6, and vapour- density determinations 
have shown that at temperatures just above the boiling-point 
many of the molecules are octatomic (S 8 ). On raising the 
temperature the vapour density gradually falls, the molecules 
dissociating into S 6 , S 4 , S 8 , and S 2 , while at 2,000 the vapour 
contains even the monatomic molecules S. The colour of the 
vapour changes in the process : at 450 it is orange-red, at 
500 deep red, and above 600 pale yellow. BILTZ has shown 
that at 650 sulphur vapour consists chiefly of S 3 molecules. 

Atomic Weight. The equivalent of sulphur was deter- 
mined by STAS, who passed sulphur vapour over a weighed 
quantity of pure silver. He found that 107-88 parts by 
weight of silver will combine with 16-035 of sulphur. The 
least weight of sulphur found in the gram-molecular weight 
of any of its compounds is 2 x 16-035 gms. Hence the 
atomic weight of the element is 2 x 16-035 = 32-07. 


Sulphuretted hydrogen, H 2 S. This compound is alter- 
natively called hydrogen sulphide or, since its aqueous 
solution has an acid reaction, hydrosulphuric acid. It is 
formed synthetically when hydrogen is passed through boiling 

H a + S = H 2 S. 

When hydrogen and sulphur vapour are heated together, 


equilibrium is set up between hydrogen, sulphur, and 
sulphuretted hydrogen 

H 2 + S ^r H 2 S. 

Since the latter substance is exothermic we should expect by 
Le Chatelier's principle that raising the temperature would 
diminish the proportion of sulphuretted hydrogen in the 
equilibrium mixture ; this is found to be the case. At 1,700 
the gas contains practically no undissociated sulphuretted 

Sulphuretted hydrogen occurs naturally in certain volcanic 
gases, and it is possibly by the incomplete combustion of this 
substance that the sulphur deposits in volcanic regions have 
been formed 

2H 2 S + 2 = 2H 2 O + S 
or 2H 2 S + 3O 2 = 2S0 2 + 3H 2 

followed by 2H 2 S + S0 2 = 2H 2 + 3S, since sulphuretted 
hydrogen immediately reacts with moist sulphur dioxide 
when the two gases are mixed. 

In the laboratory sulphuretted hydrogen is made by the 
action of dilute hydrochloric acid upon ferrous sulphide 
FeS + 2HC1 - FeCl 2 + H 2 S. 

A purer gas is obtained from antimony sulphide and con- 
centrated hydrochloric acid 

Sb 2 S 3 + 6HC1 - 2SbCl 3 + 3H 2 S. 

As the gas is very often required in the laboratory, forms 
of apparatus have been devised which are automatic in 
action ; sulphuretted hydrogen is produced when the tap is 
turned on and the action stopped when the tap is turned off. 
The well-known apparatus devised by KIPP is generally used. 

Properties. Sulphuretted hydrogen is a colourless gas with 
a sweetish and somewhat unpleasant odour of rotten eggs. 
It is a poisonous substance and, if inhaled, rapidly produces 
headache. It will burn in air with a blue flame, forming 
steam and sulphur dioxide if the supply of air be sufficient 

2H 2 S + 30 2 = 2H 2 O + 2S0 2 . 
If the supply of air is insufficient sulphur is deposited. 


The gas reacts with chlorine water, bromine water, or iodine 
suspended in water, yielding a precipitate of sulphur and a 
solution of halogen hydracid 

H 2 S + X 2 = 2HX + S, where X = 01, Br, or I. 

When sulphuretted hydrogen is sparked, it is decomposed 
into hydrogen and solid sulphur. The volume of the residual 
hydrogen is found to be equal to the original volume of 
sulphuretted hydrogen. Hence, by Avogadro's hypothesis, 
1 molecule of the gas contains 1 molecule of hydrogen, H 2 , 
.'. formula is HgS^.. 

The vapour density is 17, /. M. Wt. = 34. But of these 
34 parts 2 are hydrogen, .*. 32 are sulphur. 
But the atomic weight of sulphur = 32, 

.*. x = 1, and the formula is H 2 S. 

Sulphuretted hydrogen dissolves in water to the extent of 
290 c.c. in 100 c.c. of water at 20 C. The solution reacts 
feebly acid and is of great importance in analytical chemistry. 
Sulphuretted hydrogen is a dibasic acid and in solution ionizes 
in two stages 

(i) H 2 S ^= H' + HS' 

(ii) HS' ;= H* + S" 

The second stage occurs to only a very small extent. In 
the presence of an acid, even the first stage is almost com- 
pletely suppressed, whereas addition of an alkali, by removing 
hydrogen ions to form un-ionized water, will increase the 
formation of HS' and S" ions. Many metals form sulphides 
that are insoluble in water ; they can be divided into two 
groups, (a) those whose sulphides are so insoluble that even 
the small concentration of S" ions present in an acidified 
solution of sulphuretted hydrogen is large enough to cause 
their precipitation from solutions of the metallic salts, and 
(6) those whose sulphides are more soluble but still so slightly 
soluble that the concentration of S" ions in an alkaline 
-solution is sufficient to precipitate them, while that in acid 
solutions is insufficient. The former metals, lead, mercury ^ 
bismuth, copper, cadmium, arsenic, antimony and tin* 



are precipitated as sulphides in Group II of the analytical 
tables ; the latter metals include iron, cobalt, nickel, zinc 
and manganese, the last four of which are precipitated 
as sulphides in Group IV, the iron having been previously 
removed by another process in Group III. 

Hydrogen per sulphide, H 2 S 2 , is known. It is analogous to 

H 2 2 and is prepared 
as a yellow oil with a 
pungent smell by adding 
a cold solution of calcium 
polysulphide slowly to 
well-cooled hydrochloric 

0"0 ^^ Oxy -Compounds of Sul- 


' Sulphur forms four 

S 2 3 , sulphur sesqui- 

SO 2 , sulphur dioxide 
or sulphurous anhydride. 

SO 3 , sulphur trioxide 
or sulphuric anhydride. 

S 2 7 , sulphur hept oxide 
or persulphuric anhy- 

Sulphur dioxide, S0 2 . 
This gas was first pre- 
pared by PBIESTLEY in 
1774, who made it by 
the action of hot concentrated sulphuric acid upon mercury 

Hg + 2H 2 S0 4 = HgS0 4 + 2H 2 + S0 2 . 

LAVOISIER in 1777 showed that it is an oxide of sulphur. 

Preparation. Sulphur dioxide is formed * when sulphur 
burns in air or oxygen, but prepared in this way is always 
mixed with traces of solid sulphur Jn'oxide, S0 8 , which may 
cause the gas to have a cloudy appearance. Sulphur dioxide 

FIG. 102. Preparation of 
Sulphur Dioxide. 


is usually prepared by heating copper with strong sulphuric 
acid. The first reaction that occurs may be-- 

Cu + H 2 S0 4 = CuS0 4 + H 2 . 

The " nascent " hydrogen then finds two things to reduce, 
viz., the excess of sulphuric acid and the copper sulphate, 
and reduces some of each 

(i) H 2 S0 4 + H 2 = 2H 2 O + S0 2 . 
(ii) 2CuS0 4 + 9H, - Cu 2 S + 8H 2 O + H 2 S. 
Sulphur dioxide results from the reduction of the sulphuric 
acid, and black copper sulphide from the reduction of the 
copper sulphate. The mixture left in the flask is therefore 
coloured black. It contains also some unreduced copper 

Many other substances will reduce sulphuric acid if heated 
with it, e.g., carbon 

C + 2H 2 SO 4 = 2S0 2 + 2H 2 O + C0 2 . 

This process is often used for the commercial preparation 
of sulphur dioxide. 

Sulphur dioxide is generally prepared commercially by 
roasting iron pyrites, FeS 2 , in a current of air. It is formed, 
too, by the action of a dilute acid upon a sulphite, bisulphite, 
or thiosulphate 

(i) Na,S0 3 + 2HC1 = 2NaCl + HoO + SO 2 . 
(ii) NaHS0 8 + HC1 = NaCl + H 2 6 + S0 2 . 
(iii) Na 2 S 2 3 + 2HC1 = 2NaCl + H 2 + S + S0 2 . 

Properties. Sulphur dioxide is a colourless gas with the 
pungent smell and taste of " burning sulphur." It is 2-3 
times heavier than air, and may therefore be conveniently 
collected by downward displacement. It dissolves in water to 
give an acid solution containing some sulphurous acid, H 2 S0 3 

H 2 + S0 2 = H 2 S0 3 , 
The following equilibria occur in the system 

S0 2 ;= S0 2 + H 2 ;= H 2 S0 3 ;= H' + HSO,' 
Gas. Dissolved. Sulphurous 


H- + H- + so/ 


When sulphur dioxide solution is boiled, sulphur dioxide is 
driven off, and sulphurous acid has never been isolated. 

Sulphur dioxide is readily liquefied ; it fortns a colourless, 
mobile liquid which is put on the market in " syphons." 
When the valve of such a syphon is loosened, sulphur dioxide 
gas comes off, while if the syphon is inverted liquid sulphur 
dioxide can be run out. 

Sulphurous acid. This acid is known only in dilute 
solution and in the form of its salts, the sulphites. It is a 
weak acid, although stronger than carbonic (it liberates 
carbon dioxide from carbonates), and is a reducing agent 
since it is easily oxidized to sulphuric acid 

H 2 S0 3 + oxygen = H 2 S0 4 . 

Moist sulphur dioxide and sulphurous acid bleach many 
colouring matters ; the bleaching process is usually a reduc- 
tion, resulting in the dye taking up two atoms of hydrogen 
per molecule with the formation of a colourless substance 
called the leucobase of the dye. Generally the leucobase of a 
dye is very readily oxidized back again to the dye itself, so 
that sulphur dioxide bleaching is often not permanent. Straw 
hats, for example, are bleached with sulphur dioxide, in order 
that a few weeks' wear may be enough for the yellow colour to 
return, thus necessitating a new hat every season. Yet some 
people ask of what use chemistry is to the business man ! 

Sulphurous acid probably exists in two forms in its aqueous 

/OH 0. /H 

o = < ;= >< 

X OH 0^ X OH 

{Cf. hydrogen peroxide, p. 436.) 

Calcium bisulphite, Ca(HSO 3 ) 2 , is the most important salt 
of sulphurous acid. Some 190,000 tons of sulphur are annually 
converted into this salt, which is employed in the manufacture 
of paper from wood. Wood consists chiefly of two things, 
lignin (30 per cent.) and cellulose (70 per cent.). Paper is 
made from cellulose, and the calcium bisulphite is used to 
dissolve out the lignin from wood, leaving the cellulose. 


Sodium metabisulphite, Na 2 S 2 5 , a salt much used in 
photography, is made by evaporating a solution of sodium 
bisulphite, NaHS0 3 

2NaHS0 8 - Na 2 S 2 5 + H 2 0. 

Sodium bisulphite itself is used in organic chemistry for 
the separation and purification of aldehydes and ketones, 
with which it forms addition compounds - 

C 6 H 6 . C< + NaHSO, = C 6 H 5 . C(-OH 

^0 \S0 3 Na 

Benzaldehydo. " Bisulphite compound " 

of beiizalclohyde. 

The " bisulphite compounds " of aldehydes and ketones are 
usually beautifully crystalline. 

Sulphur dioxide and sulphites are used as antichlors because 
they react with chlorine and will thus remove the last traces 
of chlorine from fabrics that have been bleached with this 

S0 2 + 2H 2 + C1 2 - H 2 S0 4 + 2HCL 

With barium chloride, sulphur dioxide solution gives a 
slight precipitate of barium sulphite 

S0 2 + H 2 + BaCU - BaS0 3 + 2HC1. 

If, however, a little potassium permanganate solution is 
added first, the colour of the permanganate is discharged since 
the sulphurous acid is oxidized to sulphuric acid and the 
permanganate reduced. The solution with barium chloride 
now gives a heavy white precipitate of barium sulphate 

(i) SO 2 + H 2 + oxygen from permanganate H 2 S0 4 
(or K 2 Mn 2 8 +5S0 2 +2H 2 0-K 2 SO^-2MnS0 4 +2H 2 S0 4 ). 

(ii) H 2 S0 4 + BaCl 2 - BaS0 4 ^ + 2HC1. 

Sulphur trioxide, S0 3 . Sulphur trioxide is a white 
substance crystallizing in fine needles. It is prepared by the 
oxidation of sulphur dioxide by means of oxygen, a catalyst 
being employed to accelerate the reaction. 

When sulphur dioxide is heated with oxygen, slow com- 
bination of the two gases occurs, but the reaction is reversible 


The forward reaction is exothermic ; increase of tempera- 
ture will therefore result in a lower proportion of the trioxide 
in the equilibrium mixture. Thus at 400 the proportion of 
trioxide at equilibrium is 98-5 per cent. ; at 700 it is 60 per 
cent., and at 900 it is practically nil. 400 then would seem 
to be the highest temperature at which a good yield of the 
trioxide can be expected, but the unaided reaction is slow at 
this temperature. However, platinized asbestos acts as an 
efficient catalyst, and a practically quantitative yield can be 
obtained in this way. 

Sulphur trioxide dissolves in water with a hissing noise, 
forming sulphuric acid, of which it is therefore the anhydride 

H 2 + S0 3 = H 2 S0 4 . 

Sulphuric acid, H 2 S0 4 . It is probable that this com- 
pound was known to JABIB IBN HAYYAN (eighth century 
A.D.). It has certainly been known for several hundreds of 
years and was prepared by the alchemists by strongly heating 
green vitriol (ferrous sulphate). " The first sulphuric acid 
works was erected by WARD in 1740 at Richmond, near 
London. He heated a mixture of sulphur and saltpetre in 
iron capsules and collected the sulphuric acid vapours in glass 
vessels of 300 litres capacity, containing a little water. The 
product was concentrated by heating the glass vessels in a 
sand-bath, but great inconvenience was caused by the facility 
with which the vessels were broken. In order to avoid this 
trouble ROEBUCK and GARBETT in 1746 replaced them by 
receptacles or chambers of lead about two metres wide, in 
which a furnace in the centre produced the acid from sulphur 
and saltpetre. In 1766 this method was first introduced into 
France by HOLKER (at Rouen) ; in 1774 LA FOLLIE also passed 
a jet of steam into the lead chamber ; and in 1793 CLEMENT 
and DESORMES showed the importance of a current of air in 
the lead chamber, which facilitated the formation of sulphuric 
acid and effected a notable saving of nitre and sulphur. 
Thus, before that time only 130 kilos of sulphuric acid were 
obtained from 100 kilos of sulphur, whereas to-day this yield 
is more tha*x doubled. They had then already correctly 


interpreted the process of sulphuric acid formation, and said 
that the nitric acid was merely a means of fixing (by means 
of nitrous vapours) the oxygen of the air which transformed 
the sulphur dioxide into sulphur trioxide. . . . Until about 
the year 1835, Sicilian sulphur had always been employed 
for the manufacture of sulphuric acid, but when the govern- 
ment of Ferdinand II of Bourbon, at Naples in 1858, conceded 
the monopoly of sulphur-mining to the House of Taix-Aycard 
and Co., of Marseilles, a concession which raised the price 
from 95. Id. to 28s. per ton, all Europe endeavoured to obtain 
sulphur from other sources, especially from pyrites." l 

Sulphuric acid is the most important chemical of commerce. 
It is used directly or indirectly in practically every art and 
trade. About 10,000,000 tons are produced annually. 


1. LEAD-CHAMBER, PROCESS. The theory of this process 
is very simple, although attempts have been made to com- 
plicate it 

(i) H 2 + S0 2 + N0 2 = H 2 S0 4 + NO 
(ii) NO + oxygen from air = N0 2 . 

A hot mixture of air and sulphur dioxide, containing also 
oxides of nitrogen, has a fine spray of water projected into 
it, when sulphuric acid is formed. Nitric oxide acts as an 
oxygen-carrier, combining with atmospheric oxygen and 
handing it on to the sulphur dioxide, thus forming sulphur 
trioxide, which is immediately converted by the water present 
into sulphuric acid. The above equations show that there 
is no loss of nitric oxide in the operation ; from this point 
of view, therefore, nitric oxide may be said to act as a catalyst 
on the reaction. 

The above theory, however, appeared to be too simple 
for many chemists, who set to work to see if they could not 
produce more complicated explanations of the reactions. 
Needless to say they succeeded, but it is refreshing to find 
that, after all, the simple equations given above really represent 
the reactions that go on better than the more complicated 

1 Molinari, General and Industrial Inorganic Chemistry, p. 254 (1912). 


ones which have been suggested. If insufficient moisture 
be present in the reaction chambers, a crystalline solid separ- 
ates. This compound is called " chamber* crystals," or 
nitrososulphuric acid, HS0 4 . NO ; with water it gives sul- 
phuric acid, nitric oxide, and nitrogen peroxide 

2HS0 4 .NO + H 2 = 2H 2 SO< + NO + N0 2 . 

There is, however, no evidence to show that chamber 
crystals are an intermediate stage under normal working 
conditions. RASCHIG assumed that nitrososulphonic acid 
(HS0 3 .NO) and nitrosisulphonic acid (H 2 NS0 5 )were formed 
as intermediate products, but this theory has since been 
proved incorrect, by DIVERS and by REYNOLDS and TAYLOR. 

The sulphur dioxide required for the process is obtained 
by roasting iron pyrites, FeS 2 , in a current of air. The gas 
issuing from the pyrites burners contains about 7 per cent, 
of sulphur dioxide, 83 per cent, nitrogen, and 10 per cent, 
oxygen. A small supply of oxides of nitrogen is now fed 
into the gas, to make up for the inevitable slight losses of 
these substances which occur in working. The oxides of 
nitrogen are prepared by the catalytic oxidation of ammonia 
by means of heated platinum in the presence of oxygen ; 
they were formerly prepared by the action of concentrated 
sulphuric acid upon nitre. Nowadays, the oxides of nitrogen 
are occasionally introduced by adding nitric acid in the Glover 

The mixture of air, sulphur dioxide, nitrogen and small 
amounts of oxides of nitrogen now passes up the GLOVER 
tower, which is made of brickwork lined with lead, inside 
which again is another lining of fireproof bricks. The tower 
is packed with broken flints, and down it trickles a mixture 
of (a) dilute " chamber " acid (65 per cent. H 2 S0 4 ) and (b) 
concentrated sulphuric acid (containing dissolved oxides of 
nitrogen) from the GAY-LUSSAC tower at the other end of 
the plant. As the hot gases pass up the Glover tower they 
are cooled to the proper temperature for reaction, they take 
up the necessary oxides of nitrogen by driving these out from 
the descending acid (remember that the oxides of nitrogen 


previously introduced are only to make up for loss ; the main 
quantity of these oxides is taken up in the Glover tower), 
and they concentrate the acid which flows down the tower. 

After the gases leave the Glover tower, they pass into a 
series of lead chambers, where the main portion of the sul- 
phuric acid is formed. (Some is formed in the Glover tower.) 
A fine spray of water is blown into the chambers from the 
roof, and " chamber acid " collects on the floor, whence it 
can be drawn off. Chamber acid is never allowed to reach 
a concentration of more than 70 per cent. H 2 S0 4 , since stronger 
acid would dissolve the lead of the chambers. 

The number and size of the chambers are so arranged that 
the conversion of sulphur dioxide into sulphuric acid is prac- 
tically complete by the time the gases leave the last chamber. 
The residual gases consist chiefly of nitrogen and oxides of 
nitrogen. To recover the latter the GAY-LUSSAC tower is 
used. This is filled with broken coke, down over which a 
stream of cold concentrated sulphuric acid flows. This 
dissolves the oxides of nitrogen, probably forming a compound 
with them, while the nitrogen passes on into the chimney 
stack. The acid from the Gay-Lussac tower is then returned 
to the Glover tower, in which it loses its oxides of nitrogen 
as already described. 

The dilute sulphuric acid (65-70 per cent.) which collects 
in the chambers is concentrated, if required, by passing a 
stream of hot air over the liquid heated in silica pans, or 
(GAILLARD process) by passing the liquid in a fine spray 
down a tower (made of acid-resisting lava) through which 
a current of hot gases from a coke generator passes in the 
opposite direction. An acid of 92-97 per cent. H 2 S0 4 collects 
at the base of the tower. 100 per cent, sulphuric acid may 
be made by adding sulphur trioxide, or sulphuric acid con- 
taining dissolved sulphur trioxide, to the 92-97 per cent. acid. 

Impurities in Chamber Acid. The chief impurities in 
commercial sulphuric acid are arsenious oxide, As 2 O 3 (from 
the pyrites), lead sulphate, PbS0 4 (from the chambers), and 
oxides of nitrogen (chiefly N 2 8 ). The arsenic is removed 
by diluting the acid and passing a stream of sulphuretted 



hydrogen through, when the arsenic is precipitated as arsenioua 

As 2 8 + 3H 2 S = As 2 S 3 * + 3H 2 

Arsenic may also be removed by adding just sufficient hydro- 
chloric acid to convert the arsenious oxide present into 
arsenic trichloride, AsCl 3 , and then blowing air through the 
liquid, when the chloride is carried off. The dark colour 

Dilute Acid 

Ifot Gases 


Fio. 103. Concentration of Sulphuric Acid (Gaillard Tower). 

often shown by commercial sulphuric acid is due to the 
charring of organic impurities. 

2. CONTACT PROCESS. The reaction 2S0 2 + 2 ; ^ 2S0 3 
has already been discussed (p. 467). It should be noted in 
addition that since a contraction of volume takes place when 
sulphur trioxide is formed from sulphur dioxide and oxygen, 
increase of pressure will aid the formation of the trioxide 
(Le Chatelier). The law of mass action, again, will show us 
that to convert all the sulphur dioxide into trioxide an excess 
of oxygen must be used. To sum up, we may say that to 
get a good yield of sulphur trioxide 


(i) The temperature must be as low as conveniently pos- 
sible, since a rise in temperature favours the reverse action, 
2S0 3 ->2S0 2 + 2 . 

(ii) A catalyst is therefore necessary to make the reaction 
proceed at a reasonable rate. 

(iii) An increase of pressure would aid the formation of 

(iv) An excess of oxygen should be used, to ensure complete 
conversion of the dioxide into trioxide. 

Since sulphur trioxide forms sulphuric acid when it com- 
bines with water, if the formation of the trioxide could be 
carried out on a commercial scale we should have a new process 
for the manufacture of sulphuric acid. 

Attempts to make sulphuric acid in this way were made 
over 100 years ago at Bristol by a vinegar manufacturer 
named PHILLIPS, but a successful process was not worked 
out until the end of the nineteenth century, when a firm of 
German chemical manufacturers, the Badische Anilin und 
Soda Fabrik, solved the problem. 

At the present day, the lead-chamber process is obsolescent, 
and will no doubt be completely superseded by the " contact " 
process at no very remote date. Various modifications of 
this process are in use ; some of the most important are 
described below. 

(i) Badische Anilin und Soda Fabrik Method. Sulphur 
dioxide from pyrites burners, mixed with excess of air, is 
passed into chambers called " scrubbers " ; into these cham- 
bers jets of steam are blown, when each particle of dust in 
the gases becomes the centre of a tiny drop of water. These 
drops of water are allowed to settle and the purified gas 
passed through coke soaked in strong sulphuric acid, where 
it is dried. This preliminary purification is necessary, as 
the dust contains arsenious oxide and other substances which 
would " poison " the catalyst (i.e. stop its action). 

The pure dry gas, consisting of sulphur dioxide and excess 
of air, is now passed through a series of tubes containing the 
catalyst, which consists of platinized asbestos, that ia, asbestos 
which has been soaked in platinum chloride solution and then 



strongly heated, when the platinum chloride loses its chlorine 
and the platinum is left in a state of extremely fine division 
throughout the asbestos. In these tubes the Teaction occurs. 
The reaction is exothermic, so that when it has been started 
by external heat it proceeds automatically, the rate of flow 
of the gases through the tubes being regulated in such a way 
that the heat evolved is just sufficient to keep the tempera- 
ture at 400-450. 

The sulphur trioxide so produced is absorbed in 98 per 



FIG. 104. Diagram of the Badische Contact Process. 
A. Ring of gas burners. B. Catalyst. 

cent, sulphuric acid in a large vat, water being run in at the 
same time at such a rate that the concentration of the acid 
in the vat remains constant at 98 per cent. When the vat 
is nearly full, the water may be turned off and the concen- 
tration of acid raised to 100 per cent. If sulphur trioxide 
is still passed it dissolves in the 100 per cent, acid and 
forms fuming sulphuric acid or oleum, a solution of the trioxide 
in sulphuric acid. 

At first attempts were made to dissolve the sulphur 





mass is 



ate^ Distributing 

Catalyst 374 




^. f^^"^^ 

trioxide in water directly, but this produced a mist of sul- 
phuric acid that filled the works and proved extremely 
unpleasant to the workmen ; moreover, much of the trioxide 
was converted into a glassy modification that dissolves 
in water much more slowly than the ordinary form. 

(ii) The Schroder -Grille Process. This differs from the 
Badische process only in the form of the catalyst. 
stead of employing plat- 
inized asbestos, magnes- 
ium sulphate crystals 
(MgS0 4 .7H 2 0) are taken, 
soaked in a solution of 
platinum chloride, and 
then heated in sulphur 
dioxide, when a spongy 
mass is obtained 
with very finely 
platinum. This 
used as the catalyst. 

(iii) The Verein chemischer 
Fabriken (Mannheim) Pro- 
cess. In this process the 
contact agent or catalyst 
is burnt pyrites, which con- 
sists chiefly of ferric oxide, 
Fe 2 3 . Combination of 
the sulphur dioxide and 
oxygen is only partial 
(60 per cent.) in this 
case, and has to be completed by platinized asbestos. 

(iv) The Selden Process, developed in America, uses oxides 
of vanadium as catalyst. These are much cheaper than 
platinum and not so susceptible to " poisons." A different 
form of contact chamber is employed, of greater efficiency. 

Comparison of Lead-chamber and Contact Processes. The 
lead-chamber process is the cheaper of the two, but produces 
a less pure acid. As for many purposes sulphuric acid is 
not required to be of a high degree of purity, chamber acid 


Catalyst 470 
sssss.\ ^^\\\^wssss? 



Gases enter 
at 338" 


Fia/105. Grille Plant. 


still has a wide sale (e.g. in the manufacture of " super- 
phosphate " fertilizers) ; but if the contact acid can be 
lowered in price it will probably kill the sale fcf lead-chamber 
acid completely. At the present day, contact acid is mainly 
employed in the manufacture of chemicals, explosives and dyes 
(especially indigo), and also in certain processes with food- 
materials such as beer. 

Properties. Sulphuric acid is a colourless oily liquid 
that can be frozen to white crystals melting at 10-5. The 
purest commercial acid (known as monohydrate, i.e., the 
monohydrate of sulphur trioxide, S0 8 .H 2 O) contains about 
97 per cent. H 2 S0 4 and 3 per cent. H 2 0. The most concen- 
trated acid obtainable by distillation under atmospheric 
pressure boils at 330 and contains 98-3 per cent. H 2 S0 4 . A 
very characteristic property of sulphuric acid is its grea-t 
avidity for water ; so much heat is evolved when the 
two are mixed that in diluting the acid the latter must 
always be added to the water, and never water to the acid. 
Addition of water to concentrated sulphuric acid might 
cause the first few drops of water to be converted into 
steam, the expansion of which would scatter the acid 

Evolution of heat is a very general sign of chemical reaction ; 
we should therefore expect that when sulphuric acid is added 
to water combination of the two occurs. In point of fact two 
hydrates of sulphuric acid, H 2 S0 4 . H 2 O and H 2 S0 4 . 4H 2 O, have 
been isolated. The great avidity of sulphuric acid for water 
makes it an excellent drying agent for gases (except ammonia, 
with which it combines to form ammonium sulphate, and 
sulphuretted hydrogen, which it oxidizes to sulphur). 

Sulphuric acid will often remove the elements of water from 
substances containing them ; thus it chars sugar 

C 12 H 22 O n = 12C + 11H 2 0, 

leaving a black mass of carbon. Paper, (C 6 H 10 5 ) n , is simi- 
larly charred, while alcohol, C 2 H 6 OH, is converted into 
ethylene, C 2 H 4 

C 8 H 5 OH = C 2 H 4 + H 8 0. 


Vapour-density determinations show that the vapour of 
sulphuric acid consists chiefly of sulphur trioxide and water 

Sulphuric acid is a " strong " acid (p. 139), but only about 
half as strong as hydrochloric and nitric acids. It has, 
however, the great advantage over them of boiling at a much 
higher temperature (330) ; it is therefore able to liberate 
them (and many other acids) from the corresponding metallic 
salts, on heating. 

It is a dibasic acid and forms two series of salts, e.g., with 
sodium it forms sodium hydrogen sulphate, NaHS0 4 , and 
normal sodium sulphate, JNa 2 S0 4 . 

When passed through a red-hot silica tube sulphuric acid 
is split up into water, sulphur dioxide and oxygen. 

The constitution of sulphuric acid is probably 

H ( 

H O 

since it is formed when water acts upon sulphuryl chloride, 

S0 2 C1 2 (p. 480). 

O x /.Cl HOH /O H 

>S/' ""+" = \S/ + 2HC1. 

o^ x ;ci HOH o^ X O-H 

Sulphur heptoxide, S 2 7} was obtained in 1877 by 
BERTHELOT, by continued sparking of a mixture of sulphur 
dioxide and oxygen 

4S0 2 + 30 2 = 2S 2 7 . 

It is a white crystalline solid that dissolves in water to 
give persulphuric acid, H 2 S 2 8 , of which it is therefore the 

Persulphuric acid, H 2 S 2 8 , can be obtained as a white 
crystalline solid by the action of chlorosulphonic acid (p. 481) 
oh 100 per cent, hydrogen peroxide 

2C1.S0 2 .OH + H 2 2 = H 2 S 2 8 + 2HC1. 
It can be prepared in aqueous solution by electrolysis of fairly 


dilute (1:1) sulphuric acid, using a very small anode. The 
HSO 4 groups that are liberated at the anode combine 
together to form persulphuric acid 

2HS0 4 = H 2 S 2 8 . 

The potassium, sodium and ammonium salts may be made by 
substituting the corresponding acid sulphate for sulphuric 
acid in the electrolysis. 

Ammonium persulphate, (NH 4 ) 2 S 2 8 , is used commercially 
as a bleaching agent. If heated with concentrated nitric acid 
it yields a mixture of ozone and oxygen. 

Persulphuric acid and the persulphates are strong oxidizing 

O.S0 2 .OH 
agents. The constitution of the acid is | 

O.S0 2 .OH 

Caro's acid, S0 2 (sometimes called permonosulphuric 

acid to distinguish it from ordinary persulphuric acid, which 
is perdisulphuric acid), was discovered in 1898 by Caro. lie 
made it by the action of concentrated sulphuric acid upon 
potassium persulphate. It is a monobasic acid and a powerful 
oxidizing agent. 

Thionic Acids. These are comparatively unimportant 
acids of the general formula H^Og, where x 2, 3, 4, 5, or 
6. A mixture of some of them is formed when sulphuretted 
hydrogen is passed into a solution of sulphur dioxide. The 
product is called " WACKEKRODEK'S solution." 

Thiosulphuric acid, H 2 S 2 3 , is chiefly of importance in 
the form of its sodium salt, sodium thiosulphate, Na 2 S 2 3 . 5H 2 
the " hypo " of photographers. This is made by boiling a 
solution of sodium sulphite with sulphur and is a white cry- 
stalline solid. It is used in volumetric analysis for the 
estimation of iodine, with which it reacts to form sodium 
iodide and sodium tetrathionate 

2Na 2 S 2 8 + I 2 = 2NaI + Na 2 S 4 6 . 

It is also used to " fix " photographic plates and prints by 


dissolving out unchanged silver halides, with which it forms 
soluble " mixed " thiosulphates 

AgBr + Na 2 S 2 3 = NaAgS 2 8 + NaBr. 

Sodium thiosulphate will react with chlorine 

Na 2 S 2 8 + 4C1 2 + 5H 2 - Na 2 S0 4 + H 2 S0 4 + 8HC1, 
and hence is employed as an antichlor for fabrics after chlorine- 
bleaching. The reaction with bromine is similar (contrast 
action with iodine, above). 

It was formerly manufactured from the alkali -waste of the 
Leblanc soda process (p. 201). On exposure of the waste 
(CaS) to air in presence of water calcium thiosulphate is 

2CaS + 20 2 + H 2 = CaS 2 3 + Ca(OH) 2 . 

On addition of sodium carbonate calcium carbonate is 
precipitated and a solution of sodium thiosulphate left 
. CaS 2 8 + Na 2 C0 3 = CaC0 3 + + Na 2 S 2 3 . 

As the supplies of alkali- waste will soon be exhausted, 
alternative methods of manufacture have been initiated. 
These include (i) the action of sulphur dioxide upon a solution 
of sodium carbonate containing suspended sulphur 

H 2 S0 3 + Na 2 C0 3 = Na 2 S0 3 + H 2 + CO, 
Na 2 S0 3 + S - Na 2 S 2 3 

and (ii) the action of sulphur dioxide upon a mixture of 
sodium bisulphite and sodium sulphide 

2NaHSO 3 + 2Na 2 S + 2S0 2 = 3Na 2 S 2 3 + H 2 0. 
Hydrosulphurous acid, H 2 S 2 4 , is important commer- 
cially in the form of its sodium salt, Na a S 2 4 , which is exten- 
sively employed in dyeing with indigo. It is made by the 
action of zinc upon a solution of sodium bisulphite, and is a 
white crystalline solid with powerful reducing properties. 

Zn + 4NaHSO 8 = ZnS0 8 + Na 2 S0 3 + 2H 2 + Na 2 S 2 4 . 

Sulphur sesquioxide, S 2 8 , is a blue solid made by the 
action of flowers of sulphur upon molten sulphur trioxide. It 
is decomposed by water into sulphuric acid, thiosulphurio 
acid and sulphur. 


Halogen Compounds of Sulphur. Sulphur forms the 
following compounds with the halogens. They are of little 
importance. SF 6 , sulphur hexafluoride, % colourless gas ; 
S 2 C1 2 , sulphur monochloride, yellow liquid ; SC1 4 , sulphur 
tetrachloride, reddish brown liquid ; SaBr^ sulphur mono- 
bromide, red liquid. 

The oxyhalogen compounds of sulphur are more important. 
The chief are thionyl chloride, SOC1 2 ; sulphury 1 chloride, 
SO 2 C1 2 ; and chlorosulphonic acid, HO.SO 2 .C1. 

Thionyl chloride, SOC1 2 , is a colourless fuming liquid 
made by the action of sulphur trioxide upon sulphur mono- 

S0 3 + S 2 C1 2 = SOC1 2 + S + S0 2 , 

or by the action of phosphorus pentachloride upon sulphur 

S0 2 + PC1 5 = SOCl a + POC1 3 . 

It is decomposed by water, giving a mixture of sulphurous 
and hydrochloric acids 

SOC1 2 + 2H 2 O = H 2 S0 8 + 2HC1. 

The relationship between thionyl chloride and sulphurous 
acid is expressed by the following structural formulae 

/Cl /OH 

= S< O - S< 

\C1 . X OH 

Thionyl chloride. Sulphurous acid. 

Compounds derived from acids by replacement of hydroxyl 
groups of the latter by chlorine are called acid chlorides ; thus 
acetyl chloride, CH 3 .CO.C1, is the acid chloride of acetic acid, 
CH 8 .CO.OH, and benzoyl chloride,, C 6 H 6 .CO.C1, is similarly 
related to benzoic acid, C 6 H 5 .CO.OH. 

Thionyl chloride is therefore the acid chloride of sulphurou * 

Sulphury 1 chloride, SO a Cl a , is the acid chloride of sul- 
phuric acid 

OH Qv /Cl 



Sulphuric acid. Sulphury! chloride. 

* SULPHUR 481 

It can be made synthetically by the action of chlorine upon 
sulphur dioxide, preferably in the presence of camphor as a 

S0 2 + C1 2 = SO 2 C1 2 . 

It is also formed when phosphorus pentachloride, PC1 6 , acts 
upon concentrated sulphuric acid 

H 2 S0 4 + 2PC1 5 = S0 2 C1 2 + 2POC1 3 + 2HC1. 

Sulphuryl chloride is a colourless fuming liquid, the struc- 
ture of which is sufficiently proved by its synthesis. It is 
decomposed by water, forming sulphuric and hydrochloric 

SO 2 C1 2 + 2H 2 O = H 2 S0 4 + 2HC1. 

Chlorosulphonic acid, HO.S0 2 .C1, is intermediate in 
structure between sulphuric acid and sulphuryl chloride 

XOH O, AK (X ,Cl 

>\ 7 S \ 

OH cr Na cr \ci 

Sulphuric acid. Chlorosulphonic acid. Sulphuryl chloride. 
It can be made synthetically by the action of hydrochloric 
acid gas on sulphur trioxide 

S0 8 + HC1- HO.SO 2 .C1, 

and in other ways. It is a colourless fuming liquid that 
reacts with water to form sulphuric and hydrochloric acids 

HO.S0 2 .C1 + H 2 O - H 2 S0 4 + HC1. 

Chlorosulphonic acid is used to some extent in the manufac- 
ture of dyes. It is prepared commercially by passing hydro- 
chloric acid gas into oleum (p. 474). 


1. What is the origin of the word sulphur ? Describe the properties 
of the element. 

2. Give an account of the extraction of sulphur. 

3. Is the classification of oxygen with sulphur justified by the chemi- 
cal behaviour of these two elements ? 

4. Discuss the phenomenon of allotropy, with special reference to 
the allotropes of sulphur and oxygen. 

6. Explain the use of sulphuretted hydrogen in qualitative analysis. 



6. Contrast the bleaching action of chlorine with that of sulphur 

7. How is sulphuric acid manufactured ? 

8. What do you know of the oxyacids of sulphur, excluding sulphuric 
and sulphurous acids ? 

9. Show, by means of formulae, the relationship between sulphuryl 
chloride, chlorosulphonio acid, thionyl chloride, sulphurous acid and 
sulphuric acid. 


GROUP VI, Sub-group A 


In chemical properties these elements show many resem- 
blances to one another, but scarcely any to the other members 
of the group. They are similar in many respects to iron, 
cobalt, and nickel. 


Group in Periodic System : VI ; Symbol : Cr ; Valency : 
2, 3, or 6; Atomic Weight: 52-0; Melting-point: 1,700; 
Specific Gravity: 6-8-7-1. 

History. A mineral called crocoisite was discovered by 
LEHMANN, in 1762. VAUQUELIN and KLAPROTH, in 1797, 
showed it to be the lead salt of a peculiar acid. From this 
acid Vauquelin succeeded in isolating an impure specimen 
of a new metallic element which was later called chromium 
on account of its characteristic property of forming coloured 
compounds. In 1857 DEVILLE prepared the metal in a state 
of approximate purity, but it was not until 1894 that chemi- 
cally pure chromium was made, by MOISSAN. 

Occurrence. The chief ore of chromium is chromite or 
chrome iron ore, FeO.Cr 2 8 . Crocoisite, PbCr0 4 , occurs in 
smaller quantities and is rarer. 

Preparation. Chromium of a high degree of purity may 
easily be made by the thermite process. Chromium sesqui- 
oxide, Cr a 3 , is mixed with a slight excess of aluminium 
powder in a fireclay crucible, and the mixture fired by igniting 



a small heap of powdered potassium chlorate and magnesium 
powder placed on the top. A very intense heat is produced, 
and the chromium oxide is reduced to chromium, which is 
found after the reaction as a mass of metal at the bottom of 
the crucible 

Cr 2 O 8 + 2A1 = 2Cr + A1 2 3 . 

Properties. Chromium is a greyish metal with a silvery 
lustre. Its specific gravity is 6-8, melting-point about 1,700, 
and boiling-point 2,200. It is stable in moist air at ordinary 
temperatures, biit burns with a bright flame when strongly 
heated. It is soluble in dilute hydrochloric and sulphuric 
acids, even in the cold, forming chromous salts ; concentrated 
nitric acid renders it passive (p. 551), probably owing to the 
formation of a thin coherent and protective coating of 

Alloys of chromium with iron, and with iron and nickel, are 
used industrially for armour-plating and many other purposes. 
" Chrome steel " is very hard and exceedingly tough, and is 
extremely resistant to acids. The so-called " stainless st^el " 
is not a true steel, but an alloy of chromium, iron, and some- 
times molybdenum containing about 12-15 per cent, of 
chromium. Chromium-plating is carried out by the electrolytic 
deposition of chromium from a solution of chromic acid and 
chromium sulphate (p. 486) to which chromium carbonate 
has been added. Useful as it is, there are indications that 
it may be replaced by cadmium plating. Chromium com- 
pounds (e.g. chrome alum and sodium dichromate) are used 
in tanning leather, in the manufacture of dyes and pigments, 
and in the photographic industry. 

COMPOUNDS OF CHROMIUM. Chromium forms three oxides, 
two of which are basic and one acidic : 
Chromous oxide, CrO. Basic. 
Chromium sesquioxide, Cr 2 O 3 . Basic. 
Chromium trioxide, CrO 8 . Acidic. 

Chromous oxide, CrO, is obtained as a black powder 
when chromium amalgam is allowed to oxidize in the air. 

Chromous hydroxide, Cr(OH) 2 , is a yellow powder 
precipitated by addition of caustic soda to a solution of 


chromous chloride. It is unstable and readily oxidizes into 
chronuc hydroxide 

2Cr(OH) 2 + H 2 O + oxygen = 2Cr(OH) 8 . 

Chromous chloride, CrCl a , is formed when chromium 
dissolves in dilute hydrochloric acid, but is usually prepared 
by the reduction of chromic chloride with nascent hydrogen, 
e.g., with zinc and hydrochloric acid. 

2CrCl 3 + H 2 (nascent) = 2CrCl a + 2HC1. 

The anhydrous salt is conveniently made by heating 
chromic chloride in a current of hydrogen ; it is a white 
crystalline solid. 

Aqueous solutions of chromous salts, and the hydrated 
salts themselves, are blue ; the anhydrous salts vary in 
colour, for while chromous chloride is white, chromous acetate 
is red. 

Chromium sesquioxide or chromic oxide, Cr 2 8 , can 
be made in a number of ways, the easiest being to heat 
ammonium dichromate 

(NH 4 ) 2 CrA = Cr 2 8 + 4H 2 + N 2 . 

Obtained in this way it is a voluminous green powder. It is 
used as a paint (" chrome-green ") and also for making coloured 
glass, to which it imparts a fine green colour. 

Chromic hydroxide, Cr(OH) 3 , is obtained as a bluish- 
green gelatinous precipitate on addition of caustic alkali to 
a solution of a chromic salt. Under suitable conditions it 
will dissolve in excess of caustic soda or potash to form salts, 
the chromites ; it will also dissolve in dilute acids to form 

chromic salts, and is therefore both basic and acidic. 


Chromic chloride, CrCl 3 , is prepared anhydrous by heat- 
ing a mixture of chromium sesquioxide and carbon in a 
current of chlorine 

Cr 2 8 + 30 + 3d 2 = 2Cr01 8 + SCO. 

The chromic chloride sublimes over in pale violet crystalline 
scales. A remarkable property of anhydrous chromic chloride 
is that it will not dissolve in water, even on heating, unless a 


trace of chromous chloride is present, when it dissolves easily, 
forming a green solution from which green crystals of the 
hexahydrate CrCl 3 .6H 2 may be obtained. ^ 

Chromic chloride hexahydrate exists in three isomerio 
forms. In aqueous solution one form yields three chlorine 
ions per molecule, the second gives two, and the third only one. 
To account for this behaviour, WERNER assumes that some 
or all of the water of crystallization forms an integral part of 
the molecule and is not split off in solution ; he then represents 
the first form as the anhydrous chloride of the ter^alent group 
[Cr(H 2 0)], the second as the monohydrated chloride of the 
bivalent group [Cr(H 2 0) 6 Cl], and the third as the dihydrated 
chloride of the univalent group [Cr(H 2 0) 4 Cl 2 ], or 
(i) [Cr(H 2 0) 6 ]Cl 3 . 
(ii) [Cr(H 2 O) 6 Cl]Cl 2 .H 2 O. 
(iii) [Cr(H 2 0) 4 Cl 2 ]C1.2H 2 0. 

For further information on this point, and on Werner's 
Theory of Valency in general, the student is referred to 
larger textbooks. 

Chromic sulphate, Cr 2 (S0 4 ) 8 , may be made by dissolving 
the hydroxide in sulphuric acid. It forms violet crystals 
containing 16 or 18 molecules of water of crystallization 
according to the method of extraction. It is most often met 
with in the form of its double salt with potassium sulphate, 
chrome alum, K 2 S0 4 .Cr 2 (S0 4 ) 8 .24H 2 0. Chrome alum is a 
dark purple crystalline substance that is usually prepared 
by passing sulphur dioxide into a solution of potassium 
dichromate acidified with sulphuric acid 

K 2 Cr 2 7 + 3S0 2 + H 2 S0 4 = K 2 S0 4 + Cr 2 (S0 4 ) 8 + H 2 0. 
On evaporation of the solution, chrome alum separates. It 
is isomorphous with alum, K 2 SO 4 .A1 2 (S0 4 ) 3 .24H 2 O, and 
is a double salt, i.e., in solution it gives all the ions which 
would be given by its constituent salts separately. (Contrast 
a typical complex salt, such as potassium ferrocyanide, 
K 4 Fe(CN) 6 , which gives none of the ionic reactions of ferrous 

Chrome alum has become of considerable importance of 


late years in the tanning of leather. Leather which has been 
soaked in chrome alum solution (" chrome-leather ") is much 
more durable than leather tanned in the ordinary way ; it 
is sstid that the boots of Chicago policemen needed repairing 
three times as seldom when they were shod with chrome 
leather as when they wore the ordinary kind. The author, 
however, is uncertain whether there are any policemen in 

Crystals of chrome alum are often found in the cells of a 
" Bichromate " battery that needs recharging. 

Chromium trioxide, Cr0 8 , separates in the form of 
delightful red needles on adding concentrated sulphuric acid 
to a concentrated solution of sodium or potassium dichromate 
and allowing the liquid to stand in a cool place. The crystals 
are deliquescent and dissolve in water to give a strongly acid 
solution containing dichromic acid 

H 2 + 2Cr0 3 > H 2 Cr 2 T 

One might have expected to get chromic acid, H 2 Cr0 4 (especi- 
ally as salts of this acid are well known) 

H a O + CrO, = H 2 Cr0 4 ; 

cryoscopic and conductivity experiments on the solution, 
however, show that it is the dichromic acid which is formed. 
When heated to 250, chromium trioxide splits up into 
chromic oxide and oxygen 

4Cr0 3 - 2Cr 2 8 + 30 2 , 

and the readiness with which this oxygen is lost makes 
chromium trioxide a strong oxidizing agent. Thus, a sus- 
pension of it in acid must be filtered through glass wool, since 
it chars paper, while if alcohol is dropped on the solid trioxide 
the mixture inflames. 

Chromates and Dichromates. These compounds are 
salts of two different acids, neither of which has been isolated, 
and one of which (chromic acid) is unknown even in solution 
The chromates are salts of chromic acid, H 2 Cr0 4 , and the 
dichromates are salts of dichromic acid, H 2 Cr 2 O 7 . It will be 
at once apparent that it is incorrect to call the potassium salt 


of dichromic acid potassium 6i-chromate, as is so often done ; 
potassium 6ichromate does not exist, but if it did it would 
have the formula KHCr0 4 , corresponding* to potassium 
bisulphate, KHS0 4 . 

Potassium chromate, K 2 Cr0 4 , may be made by neutral- 
izing a solution of dichromic acid or potassium dichromate 
with potash 

K 2 Cr 2 7 + 2KOH = 2K 2 Cr0 4 + H 2 0. 

On evaporation of the solution the potassium chromate 
separates in the form of yellow crystals. In solution, potas- 
sium chromate is used as a reagent in qualitative analysis, 
since it gives precipitates of insoluble chromates with 
many metals, e.g., silver (brick red), lead (yellow), barium 

Potassium dichromate, K 2 Cr 2 7 , is manufactured from 
chrome iron ore by fusing the ore in a reverberatory fur- 
nace with sodium carbonate and lime with free access of 

4FeO.Cr 2 8 + 8Na 2 C0 3 + 8CaO + 70 2 

= 2Fe 2 3 + 8Na 2 Cr0 4 + 8CaC0 8 . 

The fused mass, after cooling, is powdered and treated with 
a warm dilute solution of sodium carbonate (to decompose 
any calcium chromate present), the insoluble matter allowed 
to settle, and the clear liquid just acidified with sulphuric 
acid and evaporated to crystallization. This yields the 
sodium salt, Na 2 Cr 2 7 .2H 2 0. 

2Na 2 CK) 4 + H 2 S0 4 = Na 2 S0 4 + Na 2 Cr 2 7 + H 2 O. 


To obtain the potassium salt, the solution of sodium 
dichromate is mixed, before evaporation, with potassium 
chloride ; the mixture of salt and potassium dichromate so 
formed can be separated by fractional crystallization. Elec- 
trolytic processes are now coming into use. 

Potassium dichromate forms orange-red anhydrous crys- 
tals. It dissolves in water, and the acidified solution is used 
as an oxidizing agent, especially in organic chemistry and 


in volumetric analysis. In the presence of sulphuric acid 
the oxidizing action may be represented by the equation 

K 2 Cr 2 7 + 4H 2 SO< = Cr 2 (S0 4 ) 8 + K 2 S0 4 + 4H 2 + 30. 

Ferrous salts may be estimated in solution by titration with 
dichromate in presence of sulphuric acid, using potassium 
ferricyanide as external indicator. 

Sodium dichromate is much cheaper and much more soluble 
than the potassium salt, and is equally good for ordinary 
oxidations, e.g., preparation of acetaldehyde. It is not so 
convenient for volumetric analysis, as it is deliquescent and 
therefore cannot be weighed out directly. 

All chromates and dichromates are excessively poisonous, 
owing to the fact that they destroy the red corpuscles of the 

Distillation of a mixture of potassium dichromate, salt, and 
concentrated sulphuric acid yields dark red vapours which 
condense to a red liquid, called chromyl chloride, Cr0 2 Cl|. 
This substance may be regarded as the " acid chloride " of 
chromic acid 

Chromic acid, ^>C r \ ; Chromyl chloride, 

It boils at 117 and readily reacts with water, giving hydro- 
chloric and dichromic acids. With a concentrated solution of 
potassium chloride it reacts to form the red crystalline com- 
pound called PEIAGOT'S salt or potassium chlorochr ornate, 


Cr0 2 


7 ci 

Cr0 2 Cl 2 + KC1 + H 2 = CrO 2 + 2HC1. 


Perchromic acid, HCrO 5 (?) On addition of hydrogen 
peroxide to a solution of potassium dichromate acidified with 
sulphuric acid, a deep blue compound is produced that 
rapidly decomposes in aqueous solution but is much more 



stable in ether. This blue compound is called " perchromic 
acid " and was considered by MOISSAN (who isolated certain 
of its salts) to have the composition HCr0 5 ; it" may, however, 
be H 8 Cr0 8 , since crystals of this substance, with two mole- 
cules of water of crystallization, have been prepared by the 
action of nearly pure hydrogen peroxide upon chromium 
trioxide dissolved in an inert organic solvent at low tempera- 


1. Exercise your chemical ingenuity by justifying the grouping of 
chromium with oxygen and sulphur. 

2. Describe the preparation and properties of potassium dichromate. 

3. How would you make (a) chromyl chloride, (6) chromium trioxide, 
from potassium dichromate ? 

4. Write an account of the preparation and properties of chromic 


TYPICAL ELEMENTS : Fluorine, Chlorine. 
Sub-group A : Manganese. 

Sub-group B (similar to typical elements) : Bromine, 


The " halogen " elements, fluorine, chlorine, bromine and 
iodine, form a very well-marked family of elements, closely 
similar in chemical and physical properties. Manganese, 
however, would never be classed with the halogens on the 
ground of chemical relationship. 

The student will find it a useful exercise to tabulate the 
resemblances of the halogens to one another ; he may also 
tax his ingenuity in finding points of similarity between the 
halogens and manganese, other than the fact that chlorine 
forms an oxide C1 2 7 and manganese an oxide Mn 2 7 . 

It may be well to point out here that examiners are very 
fond of the question, " Why are the elements (this, that 
and the other) classified together ? " since they can thus 
test the candidate's knowledge of the chemistry of several 
elements in an economical way. The most popular elements 
for this purpose are the alkali-metals, halogens, and nitrogen 


Oroup in Periodic System : VII ; Symbol : F ; Valency : 
I j Atomicity : 2 ; Atomic Weight : 19-0. 



History. In the form of its compound with calcium, 
CaF 2 (calcium fluoride, or fluorspar), fluorine has been known 
since the days of " BASIL VALENTINE " (about 1600). 
SCHEELE was the first (1771) to prepare hydrofluoric acid, 
which was more fully investigated by GAY-LUSSAO and 
THENABD. Fluorine was *first prepared in the elementary 
state by MOISSAN in 1886. 

Occurrence. In common with the other halogen ele- 
ments, fluorine is too active to be found in nature in the 
uncombined state. It occurs chiefly as fluorspar, CaF 2 , 
which is found in many parts of the world, especially in 
Derbyshire and Mexico ; as cryolite or " ice-stone " (Na 3 AlF 6 ), 
which occurs in quantity in Greenland ; and in fluor-apatite 
(CaF 2 .3Ca 3 (PO 4 ) 2 ). In small quantities it is very widely 
distributed in sea-water, the enamel of teeth, oyster- shells, etc. 
Preparation. The preparation of fluorine proved to be 
a very difficult task. Methods analogous to those used in 
the preparation of chlorine (q.v.) proved useless, owing to the 
fact that the fluorine, even if liberated, was so active that it 
immediately entered into combination again with some of the 
other substances present. The French chemist Moissan, fol- 
lowing DAVY, tried to prepare it by electrolysis of an aqueous 
solution of hydrofluoric acid, HF, but found that ozonized 
oxygen, and not fluorine, was obtained at the anode. He 
therefore tried again with pure anhydrous hydrofluoric acid, 
kept liquid by being surrounded with a freezing-mixture. 
This method was also unsuccessful, since anhydrous liquid 
hydrofluoric acid proved to be a non-conductor. To over- 
come this difficulty, Moissan at length hit on the happy idea 
of dissolving some potassium hydrogen fluoride (KHF 2 ) in 
the liquid hydrofluoric acid, thus rendering the liquid a 
conductor. He found that on electrolysis this solution gave 
at the anode a pale yellow gas which proved to be fluorine. 
The apparatus Moissan employed was a platinum U-tube, 
fitted with side delivery-tubes and closed by stoppers made 
of transparent fluor-spar. Into this U-tube dipped the 
electrodes, which were made of an alloy of platinum and 
iridium. The tube was of about 160 c.c. capacity, arid in 



it were placed 100 c.c. of anhydrous liquid hydrofluoric acid 
and 20 grams of potassium hydrogen fluoride. On electrolysis 
by means of a current of 15 amperes and 50 volts, Moissan 
obtained fluorine at the rate of 5 litres per hour. The U-tube 
was kept at a temperature of 23 C. by immersion in a 
bath of boiling liquid methyl chloride. [See Fig. 106.] 

Moissan afterwards found that a copper tube (but not 
copper electrodes) could be used in the preparation of fluorine, 
since the copper soon becomes covered by a coherent coating 
of copper fluoride, CuF 2 , which prevents further action. 

To show that the gas 
he obtained was fluo- 
rine, and nothing else, 
Moissan absorbed a 
certain weight of it in 
sodium, and weighed 
the sodium fluoride 

Properties . Fluo- 
rine is a pale greenish- 
yellow gas of density 
19 (H = 1), corres- 
ponding to the formula 
F 2 . Liquid fluorine, 
first obtained by Mois- 
san and Dewar in 1897, 
is a transparent yellow liquid boiling at 187 C., and freezing 
at 233 C. to a pale yellow solid. Fluorine is one of the 
most chemically active elements known ; it combines with 
practically all other elements except nitrogen, chlorine, and 
the inactive gases, forming fluorides. Metals, especially 
when heated, take fire in fluorine ; so do sulphur, phosphorus 
and even carbon in a finely divided state. Fluorine and 
hydrogen combine with explosion even in the dark, forming 
hydrofluoric acid. Water is decomposed by fluorine, hydro- 
fluoric acid and ozonized oxygen resulting. Perfectly dry 
fluorine attacks glass only extremely slowly, if at all, but if 
even a trace of moisture be present rapid action takes 

FIG. 106. Preparation of Fluorine. 


place, and silicon tetrafluoride (SiF 4 ) and water are formed. 
The atomic weight of fluorine has been determined in 
many ways, e.g., by the conversion of sodium fluoride into 
sodium sulphate, and by the formation of sodium chloride 
from sodium fluoride by passing a current of dry hydrochloric 
acid gas over the latter substance. The value adopted is 
F = 19-0. 


Hydrofluoric acid or Hydrogen fluoride, HF. 

Preparation. This substance may be obtained by the 
direct union of its elements and in other ways, but it is gener- 
ally prepared by the action of strong sulphuric acid upon 
calcium fluoride 

CaF 2 + H 2 S0 4 = CaS0 4 + 2HF. 

The operation is usually carried out in a lead retort, and 
the acid which comes off is collected in a cooled receiver, or, 
if an aqueous solution is desired, the gas may be passed into 
water. The pure anhydrous acid is prepared from the aqueous 
solution as follows : the solution is divided into two equal 
parts, one of which is then neutralized with caustic potash 
solution. This gives a solution of potassium fluoride, KF 

HF + KOH = KF + H 2 0. 

To this solution the other half of the original hydrofluoric 
acid solution is added, when 

KF + HF = KHF 2 , or KF.HF 

Potassium hydrogen fluoride. 

The solution of potassium hydrogen fluoride is evaporated 
until the crystals of the salt separate out ; these are then 
carefully dried in a vacuum desiccator and the dry salt placed 
in a platinum retort connected to a platinum condenser and 
receiver, the latter being placed in a freezing mixture. On 
heating, the acid fluoride splits up into potassium fluoride 
and hydrofluoric acid 

KHF 2 = KF + HF, 

and the latter is condensed to a colourless, mobile, volatile 
liquid which collects in the receiver. It must be kept in 


strong platinum bottles, the stoppers of which can be firmly 

Properties. Hydrofluoric acid is usually met with as a 
gas, which condenses at 19 to a colourless fuming mobile 
liquid. It is readily soluble in water, forming a strongly 
acid solution. The anhydrous acid is not so active as the 
moist acid ; thus perfectly dry hydrofluoric acid has no 
action upon glass, whilst in the presence of moisture (even 
in minute quantity) rapid action takes place and the glass 
is attacked. It is this property of moist hydrofluoric acid 
which is made use of for the etching of glass. The glass 
article to be etched is covered with wax and the desired 
pattern then drawn on the wax with a sharp instrument so 
that the glass is exposed in those places where etching is 
required. It is then placed over a trough containing aqueous 
hydrofluoric acid, or a mixture of calcium fluoride and con- 
centrated sulphuric acid, the hydrofluoric acid fumes attack 
the glass in the exposed parts, and the glass is etched. 

As an acid, hydrofluoric acid is not so strong as the corre- 
sponding acids of the other halogens (HC1, HBr, HI). It 
has, however, a greater action on organic material. Its 
formula has been proved by Moissan (among others) who 
measured the volumes of hydrogen and fluorine liberated 
during the electrolysis of anhydrous liquid hydrofluoric acid 
(containing a little potassium hydrogen fluoride to make it 
conduct), and found that they were equal. Hence by Avo- 
gadro's Hypothesis the formula must be (HF) n . At a tempera- 
ture of about 75 the vapour density is 10, therefore the mole- 
cular weight is 20 and n = I ; the formula of the gas here is 
HF. At lower temperatures association takes place, that is, 
two or more molecules combine to form a more complex one, 
and the mixture of these more complex molecules with the 
simple HF ones gives a higher vapour density. 

Silver fluoride differs from silver chloride, bromide and 
iodide in being soluble in water. Many fluorides will combine 
directly with hydrofluoric acid to form acid fluorides, e.g., 
KHF 2 , so that hydrofluoric acid can act as a dibasic acid, 
although it is usually monobasic. HF, of course, must be 


monobasic ; the apparent dibasicity must be attributed to 
unusual stability of the associated H 2 F 2 molecules or to the 
fact that fluorine is occasionally capable of acting as a poly- 
valent element. 


Group in Periodic System : VII ; Symbol : Cl ; Valency : 
1 ; Atomicity : 2 ; Atomic Weight : 35-46. 

History. The history of elementary chlorine dates from 
1774, in which year it was prepared by SCHBELE by the action 
of muriatic acid (HC1) upon pyrolusite (Mn0 2 ). Chlorine in 
the form of its compounds, however especially sodium 
chloride, NaCl has been known from prehistoric times. 
Aqua regia, a mixture of hydrochloric and nitric acids, was 
prepared by the Arabian chemists in the early Middle Ages, 
by heating a mixture of salt, saltpetre and alum, or sal- 
ammoniac, saltpetre, copper sulphate and clay. GLAUBER 
(1648) prepared aqua regia by heating a mixture of aqua 
fortis (nitric acid) and salt, and spiritus salis (spirit of salt, or 
hydrochloric acid solution) by heating salt with concentrated 
sulphuric acid and dissolving the fumes in water. An earlier 
preparation of spiritus salis is that of LIBAVIUS (1595), who 
made the acid by strongly heating a mixture of salt and clay. 

In 1772 PRIESTLEY collected hydrochloric acid gas by 
means of a mercury pneumatic trough, and called it marine 
acid air on account of its preparation from sea-salt ; the 
name muriatic acid was given to it for the same reason. In 
1774 Scheele heated a solution of marine acid air, or muriatic 
acid, with pyrolusite and obtained a greenish yellow gas 
which he called dephlogisticated marine acid air, as he con- 
sidered it to be muriatic acid from which phlogiston had 
been removed by the pyrolusite. BERTHOLLET (1785) showed 
that an aqueous solution of Scheele's gas on exposure to 
light yielded bubbles of oxygen and a solution of muriatic 
acid. From this and other observations LAVOISIER (1789) 
called the gas oxymuriatic acid and regarded it as a compound 
of oxygen and muriatic acid ; he supposed the latter com- 
pound to be an oxide, Mu0 8> of a new element murium, and 
Scheele's gas to be Mu0 3 . 


GAY-LUSSAC and TH&NARD (1809), from their researches 
on these substances, supported Lavoisier's theory, but also 
pointed out that the facts agreed equally well with the 
theory that oxymuriatic acid was an element. 

SIR HUMPHRY DAVY, in 1810, observed that if oxymuriatic 
gas was a compound, as maintained by Lavoisier, it should 
be possible to split it up, and made many experiments with 
this aim. He heated metallic sodium in dry oxymuriatic 
acid but obtained nothing but sodium muriate, whereas if 
the gas were a compound, oxygen or water would probably 
have been formed. From these and many other results, 
Davy concluded that Scheele's gas was an element, and 
since " oxymuriatic acid " was obviously an inappropriate 
name for an element, after consultation with other chemists 
he re-named it chlorine, from the Greek %Ao>og, greenish- 

Within the last few years, it has been shown that ordinary 
chlorine is a mixture of two isotopic elements, one of atomic 
weight 35 and the other of atomic weight 37. Neither of 
the isotopes has yet been obtained free from the other. 

Occurrence. Owing to its great chemical activity, 
chlorine does not naturally occur in the free state. The 
chlorides of sodium, potassium and magnesium are widely 
distributed in large quantities. Sodium chloride, NaCl, is 
found as rock-salt or halite (cubical crystals), in very large 
deposits in England (Cheshire), Galicia (near the town of 
Wielickza), and Germany (Stassfurt deposits). It is also 
found dissolved, in brine-springs (Droitwich, etc.) and in the sea. 
Potassium and magnesium chlorides are found in the solid 
state at Stassfurt, and in solution in sea-water. Silver chloride, 
AgCl, or horn silver, is also found native. Chlorine is essential 
to the life of both plants and animals. 

Preparation. In the laboratory chlorine is prepared 
generally by the oxidation of hydrochloric acid 

2HC1 + oxygen = H 2 + C1 2 . 

Practically any oxidizing agent will do ; the choice is 
governed by convenience and economy. Manganese dioxide* 



and potassium permanganate are commonly employed. If 
manganese dioxide is used it is mixed with concentrated 
hydrochloric acid in a round-bottomed flask* fitted with a cork 
carrying a thistle funnel (through which more acid may be 
added if required) and a delivery tube. 
On heating, chlorine is evolved 

(i) Mn0 2 + 4HC1 = MnCl< + 2H 2 (in the cold), 
(ii) MnCl 4 == MnCl, + Cl a (on heating). 

FIG. 107. Preparation of Chlorine from Manganese Dioxide. 

With potassium permanganate, action takes place in the 
cold. The permanganate is placed in a flask fitted with a 
cork carrying a dropping-funnel and delivery tube ; con- 
centrated hydrochloric acid is run in slowly from the funnel, 
and a steady stream of chlorine comes off without the appli- 
cation of heat 

2KMn0 4 + 16HC1 = 2KC1 + 2MnCl 2 + 8H 2 + 5C1 2 . 

The gas is collected by downward displacement or over 
brine ; if required dry it is passed through a Drechsel bottle 
containing strong sulphuric acid and collected by downward 



displacement. It cannot be collected over mercury, as it 
attacks this metal, with formation of mercuric chloride. 

Chlorine may also be prepared by the action of a dilute 
acid upon bleaching-powder, CaOCl 2 . 

CaOCl 2 + H 2 S0 4 = CaS0 4 + H 2 + C1 2 . 

Pure chlorine may be made by the electrolysis of fused 
silver chloride (SHBNSTONB) using a carbon anode, or by the 

Fro. 108. Preparation of Chlorine from Potassium Permanganate. 

action of heat upon carefully dried and recrystalh'zed gold 
chloride, AuCl 3 

2AuCl a = 3C1 2 + 2Au. 

Manufacture. Chlorine is important commercially, and 
is manufactured from salt or from hydrochloric acid. 

DEACON'S PROCESS. This process employs atmospheric 
oxygen to oxidize hydrochloric acid gas to chlorine. The 
forward reaction 4HC1 + 2 ;~~ ^ 2H 2 + 2C1 2 is exothermic, 
and is therefore favoured by low temperature ; but at ordinary 
temperatures the velocity of the reaction is slow. Hence 
the conditions for successful operation are 


(i) The rate of reaction must be so far increased that the 
change takes place as quickly as possible. 

(ii) The temperature must not be so high <that equilibrium 
is appreciably moved to the left in the above equation. 

(iii) As much of the hydrochloric acid as possible must be 
converted into chlorine. 

In order to realize conditions (i) and (ii) a catalyst is em- 
ployed. Many substances were tried, but the best was found 
to be cupric chloride, CuCl 2 . In order to get a large surface 
of catalyst with least expenditure, broken bricks are soaked 
in a solution of the copper chloride and then dried. In this 
way the chloride is left in a finely divided state throughout 
the porous clay of which the bricks are made. It has been 
suggested that the action of the catalyst consists in the 
alternate formation and decomposition of cuprous chloride, 

(i) 2CuCl 2 = 2CuCl + C1 2 . 
(ii) 2CuCl + 2HC1 + atmospheric oxygen = 2CuCl 2 + H 2 0. 

It should be remembered, however, that other substances, 
for* which the above kind of explanation is not admissible, 
will act as catalysts on the reaction. 

The best temperature is found to be 450-500, and in 
practice about 70 per cent, of the hydrochloric acid is oxidized 
in one passage through the apparatus ; the remaining 30 
per cent, is absorbed in water and used again. 

The hydrochloric acid gas formerly used came from the salt- 
cake furnaces in the manufacture of sodium carbonate by the 
LEBLANG process. It was absorbed in water in a tower filled 
with coke, and then again set free from the aqueous solution 
by addition to concentrated sulphuric acid. In this way 
the impurities contained in the crude gas from the furnaces 
^vere removed. 

The purified hydrochloric acid gas obtained from this or 
other sources is mixed with excess of air (condition (iii) above) 
and the mixed gases heated in pipes to about 450, after which 
they pass into another series of tubes that contain the catalyst 
and are maintained at a temperature of about 480-500 by the 


heat of the reaction between the hot gases. The issuing gas is 
passed up a tower filled with coke over which water trickles ; 
the unused hydrochloric acid is dissolved and the remaining 
gas dried by passing through concentrated sulphuric acid. 

The chlorine obtained in this way is very dilute, being 
mixed with about fifteen times its own volume of residual air 
and nitrogen ; it is therefore not suitable for compression 
but is used directly for bleaching or for conversion into 
bleaching-powder (p. 248). 

The catalyst loses its power after about ten days and has 
to be renewed. Nowadays a mixture of cupric chloride and 
common salt is used instead of pure cupric chloride ; this 
does not require renewal so often. With the disuse of the 
Leblanc process and the general adoption of the electrolytic 
process (p. 197), the Deacon process is gradually becoming 

WELDON'S PROCESS. When concentrated chlorine is re- 
quired on a small scale it is still occasionally made by a process 
first successfully worked by WELDON in 1866. Formerly the 
Weldon process was of much more importance than it is 
now ; at the present day it has been almost entirely super- 
seded by the Electrolytic process. 

Pyrolusite (mineral manganese dioxide, about 80 per cent. 
Mn0 2 ) is treated with concentrated hydrochloric acid in 
stone chambers heated by steam, when a fairly pure chlorine 
is evolved. 

Mn0 2 + 4HC1 = MnCl 2 + 2H 2 + C1 2 . 

Manganese dioxide is expensive, and therefore to obtain 
the chlorine at a reasonable cost the manganese has to be 
recovered from the chloride, or reconverted into manganese 
dioxide which may be used again. The recovery is carried 
out by mixing the manganese chloride solution with a large 
(33-40 per cent.) excess of lime and blowing air into the 
liquid at a temperature of 60. Under these conditions 
the manganese chloride is converted first into manganous 
hydroxide by the lime 

MnCl a + Ca(OH) a = Mn(OH) a + CaCl,, 


and this into manganese dioxide by the atmospheric air 
2Mn(OH) 2 + O 2 == 2Mn0 2 + 2H 2 0. 

The manganese dioxide then combines with excess of lime to 
form calcium manganite, CaMn0 3 , or CaO .Mn0 2 , which settles 
to the bottom as a black mud, Weldon mud. This is then 
run into the chlorine stills and used instead of fresh pyrolusite. 
If excess of lime is not used in the recovery process, triman- 
ganic tetroxide, Mn 8 4 , is formed, which means that only 
one-third of the manganese is reconverted into the state of 
oxidation of manganese dioxide (Mn 8 4 = 2MnO.Mn0 2 ). 

It is clear that since the Weldon mud contains lime as well 
as manganese dioxide (calcium manganite = CaO + Mn0 2 ), 
there is a serious loss of hydrochloric acid in the form of 
calcium chloride 

CaMn0 8 + 6HC1 = CaCl 2 + MnCl 2 + 3H 2 + C1 2 , 
only 33 per cent, of the hydrochloric acid being converted 
into free chlorine. 

ELECTROLYTIC PROCESSES. Most of the chlorine used 
commercially is now prepared electrolytically, in the manu- 
facture of caustic soda, p. 197. Electrolytic chlorine is very 
pure, which is a great advantage in all cases except for the 
manufacture of bleaching-powder, for which a dilute chlorine 
is best. 

Chlorine is put on the market as a liquid, contained under 
pressure in steel cylinders. To liquefy the gas it is strongly 
cooled and then compressed ; it must be carefully dried first, 
since moist chlorine attacks the steel cylinders. 

Properties. Chlorine is a greenish -yellow gas with a 
characteristic pungent and irritating smell. It is very 
poisonous, and was largely employed in the war of 1914-18 
as a weapon of offence : sometimes with boomerang effects 
upon its users. The gas-masks used against it contained a 
mixture of charcoal and potassium permanganate. 

The vapour density of chlorine is 35-5 at ordinary tem- 
peratures, corresponding to the formula C1 2 ; at higher 
temperatures the vapour density diminishes slightly, showing 
that partial dissociation has occurred (Cf. iodine, p. 530). 


Chlorine was first liquefied by FARADAY ; it forms a golden 
yellow liquid boiling at 33-6 and solidifying on further 
cooling to a yellow crystalline solid, M.P. 102. 

Chlorine combines directly with most metals, forming salts 
called chlorides (" halogen," salt producer), and also with 
many non-metals. In general, absolutely dry chlorine is much 
less active than the ordinary moist, or not specially dried, 
gas ; water seems to act as a catalyst in these, as in many 
other, reactions. 

With hydrogen, chlorine combines explosively in the 
presence of bright sunlight, and more slowly in diffused light, 
to form hydrochloric acid gas 

H 2 + C1 2 = 2HC1. 

The reaction is very exothermic. Hydrogen and chlorine 
also explode if heated to about 260. Owing to the great 
affinity of chlorine for hydrogen, many hydrocarbons, that 
is, compounds that consist of hydrogen and carbon only, 
will burn in the element, forming clouds of hydrochloric acid 
gas and depositing the carbon as soot 

C,H y + |C1 2 = t/HCl + xO. 

, 2* 

Thus a piece of filter- paper soaked in boiling turpentine 
takes fire spontaneously in chlorine, the turpentine burning 
with a reddish smoky flame 

C 10 H 16 + 8C1 2 = IOC + 16HC1. 

Similarly a burning wax taper goes on burning if put into 
chlorine, since wax is a hydrocarbon of the type C n H 2ll + 2 . 

C n H 2fl + 2 + ( + lAi = nC + (2n + 2)HC1. 
In other cases of the reaction of chlorine with hydrocarbons 
the chlorine may combine directly with the hydrocarbon to 
form an addition product, e.g. 

C 2 H 4 + Cl a = C 2 H 4 C1 2 , 

Ethyleiie. Ethylene dichloride. 

or may replace all or part of the hydrogen in the hydrocarbon 
forming substitution products, e.g. 

CH 4 + C1 8 = CH 8 C1 + HC1. 

Methane. Chloromothane, or 

methyl chloride. 


Many metals burn when heated in chlorine, and if finely- 
divided may even ignite spontaneously in the gas at ordinary 
temperatures. Powdered antimony takes fire when sprinkled 
into a jar of chlorine, forming antimony chlorides, but sodium 
has to be heated in the gas before it will burn, although slow 
combination occurs in the cold. The anhydrous chlorides 
of metals may often be conveniently made in this synthetic 
way (e.g. ferric chloride, aluminium chloride). 

Phosphorus melts and ignites spontaneously in chlorine, 
forming phosphorus pentachloride, PC1 6 , mixed with the 
trichloride, PC1 3 . 

Chlorine will combine directly with carbon monoxide, CO, 
forming phosgene gas or carbonyl chloride, COC1 2 ; and with 
sulphur dioxide, S0 2 , forming sulphuryl chloride, S0 2 C1 2 . 
A catalyst is necessary to make these reactions proceed 
quickly ; camphor, for example, acts as a suitable catalyst 
on the second change, while finely divided charcoal is generally 
used in the first. 

Chlorine is soluble in water, the solution, which has the 
colour and smell of the gas, being called chlorine water. If 
cooled to 0, chlorine water yields greenish yellow crystals 
of chlorine hydrate, C1 2 .#H 2 where x is generally said to 
be 8. It was by heating this substance in a bent sealed tube, 
one arm of which was cooled in a freezing-mixture, that 
Faraday first liquefied chlorine ; the hydrate decomposes 
on heating, liberating chlorine. 

Chlorine and chlorine-water are good bleaching agents ; 
chlorine used for bleaching must be moist, as the dry gas 
usually has no action. The bleaching in all cases is an oxida- 
tion process ; sometimes the chlorine itself combines with 
the colouring matter to form colourless compounds, but in 
general the process is due to the action of hypochlorous acid, 
HC10, which is slowly formed by the action of chlorine on 

H 2 + C1 2 = HC1 + HC10, 

and readily yields its atom of oxygen to the dye, which is 
split up thereby into colourless substances. Note that the 
active bleaching agent is not " nascent oxygen," as is so often 


stated, but hypochlorous acid ; although it is correct, of 
course, to say that the bleaching is brought about by oxidation. 
CHLORINE AND ALKALIS. When chlorine is passed into 
cold dilute caustic soda solution, sodium chloride and hypo- 
chlorite are formed 

C1 2 + 2NaOH = NaCl + NaCIO + H 2 0. 

Analogous reactions occur with caustic potash and dilute 
milk of lime 

C1 2 + 2KOH = KC1 + KC10 + H 2 0. 
2C1 2 + 2Ca(OH) 2 = CaCl 2 + Ca(C10) 2 + 2H 2 0. 

With hot concentrated alkalis or milk of lime, however, 
the chloride and chlorate of the metal are formed 

3C1 2 + GNaOH = 5NaCl + NaC10 3 + 3H 2 0. 
3C1 2 + 6KOH = 5KC1 + KC10 3 + 3H 2 0. 
6C1 2 + 6Ca(OH) 2 = 5CaCl a + Ca(C10 3 ) 2 + 6H 2 O. 


Hydrochloric acid, or hydrogen chloride, HC1. The 
history of this gas has already been given (p. 496). It is 
obtained commercially by heating common salt with con- 
centrated sulphuric acid, when 

(i) NaCl + H 2 S0 4 = NaHS0 4 + HC1, 
(ii) NaCl + NaHS0 4 = Na 2 S0 4 + HC1. 

The gases are passed up through a tower filled with coke, 
where the hydrochloric acid is absorbed in a downward-flowing 
stream of water. By running the dilute acid thus obtained 
down over the coke again it absorbs more of the gas and so 
becomes more concentrated. The strongest hydrochloric acid 
solution obtainable is about 43 per cent, by weight. The chief 
impurities in the commercial product are chlorine, ferric 
chloride (FeCl 3 ), sulphur dioxide (S0 2 ), arsenic chloride (AsCl 8 ), 
sulphuric acid (H 2 S0 4 ), and common salt (NaCl). The purifica- 
tion of hydrochloric acid is difficult, hence the much lower 
price of the " commercial " acid. 

In the laboratory, hydrochloric acid gas is prepared by the 
action of sulphuric acid upon a chloride, generally common 



salt. The reaction takes place in the cold, but is hastened, 
as usual, by heat. The products are sodium Wsulphate and 
hydrochloric acid gas 

NaCl + H 2 S0 4 = NaHS0 4 + HC1. 

The reaction 2NaCl + H 2 S0 4 = Na 2 S0 4 + 2HC1 takes place 
only at much higher temperatures. 

The apparatus used is similar to that for preparing sulphur 
dioxide (see p. 464). The gas may be dried by bubbling 
through concentrated sulphuric acid and collected by down- 
ward displacement or over mercury. If a solution of hydro- 




Retort. Method. Funnel Method. Reversed Drechsel Bottle Method, 
FIG. 109. Methods of preventing " Sucking Back." 

chloric acid is required, the gas is passed into distilled 
water, using one of the many devices for preventing " sucking 
back," which is always liable to occur in making a solution 
of a very soluble gas. 

Properties. Hydrochloric acid is a colourless gas with a 
pungent smell. It fumes strongly in moist air, owing to the 
formation of tiny drops of hydrochloric acid solution. It can 
be liquefied by cold and pressure to a colourless liquid boiling 
at 83 and solidifying to a white crystalline solid at, 115 ; 
it was first liquefied by FARADAY in 1846. It is very soluble 
in water, 1 volume of which will dissolve 450 volumes of 
the gas at 15 C. The solution possesses strongly acid pro- 




perties, in contrast to the anhydrous liquid acid just men- 
tioned, which shows scarcely any acid reactions it will not 
act upon metals nor even upon carbonates. The solution, 
however, readily dissolves many metals with evolution of 
hydrogen, e.g. 

Zn + 2HC1 = ZnCl a + H a , 

the other products being chlorides. 

Hydrochloric acid is highly ionized in solution, and is, in 
fact, the strongest acid known, being about as strong as nitric 
acid and about one and a half 
times as strong as sulphuric. 
(See strengths of acids, p. 139.) 
The most concentrated solu- 
tion of hydrochloric acid 
obtainable at 15 C. has a 
specific gravity of 1*212, and 
contains 43 per cent, by 
weight of HC1. On distilling 
this solution hydrochloric 
acid is given off and the 
solution becomes weaker, 
the temperature gradually 
rising to 110, when no 
further change takes place 
and the solution distils over 


unchanged. The composi- 
tion of this constant-boiling solution of hydrochloric acid is 
20-24 per cent. HC1 by weight. If a more dilute solution be 
distilled, water comes off first, the temperature of the boiling- 
point gradually rising till it reaches 110, when the constant- 
boiling solution is obtained, which distils over unchanged. 
These results are shown in the figure (Fig. 110). 

This constant-boiling solution was at one time considered 
to be a compound, since its percentage composition (20'24 
per cent. HC1) is approximately that required by the formula 
HC1.8H 2 0, but it was shown that the composition of the 
constant boiling solution varies with the pressure, whereas, 


of course, the composition of a compound is independent of 
the physical conditions. Hence the constant-boiling solution 
is a mixture and not a compound. 

CHLORIDES. Most metallic chlorides are soluble in water, 
the chief exceptions being those of silver, mercury (mercunwa 
chloride, Hg 2 Cl 2 ; mercuric chloride, HgCl 2 , is soluble) and 
lead (slightly soluble) ; these metals are therefore precipitated 
as their chlorides when hydrochloric acid is added to a solution 
containing them (Group I, qualitative analysis). Metallic 
chlorides in general are not hydrolysed by water (exceptions 
antimony and bismuth chlorides, etc.), and ionize readily 
in aqueous solution (exceptions mercuric and cadmium 
chlorides). On heating with sulphuric acid they yield hydro- 
chloric acid gas ; with a mixture of manganese dioxide and 
sulphuric acid chlorine is given off on application of heat 

2NaCl + MnO a + 3H 2 S0 4 

= 2NaHS0 4 + MnS0 4 + 2H 2 + 01,. 

(This reaction is sometimes used for the preparation of 

The chlorides of non-metals are generally colourless liquids, 
for example, phosphorus trichloride, carbon tetrachloride, 
nitrogen trichloride. They are often, but not always, hydro- 
lysed when added to water ; thus phosphorus trichloride 
yields phosphorous acid and hydrochloric acid 

PC1 3 + 3H 2 = H 3 P0 3 + 3HC1. 

A rough test for deciding whether an element is a metal or 
non-metal is therefore to find the effect of water upon its 

Metal chloride + water = ionized solution, no hydrolysis. 
Non-metal chloride + water > hydrolysis, with formation 
of hydrochloric acid. 

(Note, however, the exceptions mentioned above.) 
Certain chlorides on heating lose chlorine, being reduced to 
a lower oJ$&$^ @x suffering complete decomposition into their 
elemealii^i!lfcii(T>latinum chloride, PtCl 4 , yields platinum 


and chlorine, Pt + 2C1 2 ; cuprio chloride, CuClj, gives 
cuprous chloride and chlorine 

phosphorus pentachloride gives the trichloride and chlorine 
PC1 5 ;= PC1 8 + C1 2 , 

and nitrogen chloride gives nitrogen and chlorine with ex- 

2NC1 3 = N 2 + 3C1 2 . 

Other chlorides yield the corresponding oxides if heated in a 
current of air or oxygen ; thus ferric chloride goes to ferrio 

4FeCl 3 + 30 2 = 2Fe 2 8 + 6C1 2 . 

Practically all metallic chlorides are reduced to metal if 
heated strongly in a current of hydrogen. 

Constitution of hydrochloric acid. (i) When a mix- 
ture of 1 volume of hydrogen and 1 volume of chlorine is 
exploded, the hydrochloric acid gas formed on cooling is 
found to occupy the same volume as the original mixture, 
namely 2 volumes. Hence, by Avogadro's Hypothesis, two 
molecules of hydrochloric acid gas contain 1 molecule of 
hydrogen and 1 of chlorine, .*. 1 molecule of hydrochloric acid 
gas contains \ molecule of hydrogen and \ molecule of chlorine. 
If we may assume that the molecules of hydrogen and chlorine 
are diatomic, the hydrochloric acid molecule will be HC1. 

(ii) This is confirmed by a vapour density determination, 
which gives a V.D. of 18-25 and /. a M.W. of 36-5. The 
atomic weight of chlorine is 35-5, /. the hydrochloric acid 
molecule is HC1. 

(iii) By shaking up a known volume of hydrochloric acid 
gas with sodium amalgam the acid is decomposed, yielding 
hydrogen (gas) and sodium chloride (solid) (Fig. 111). The 
volume of the residual hydrogen is found to be exactly half 
that of the hydrochloric acid gas taken, /. by Avogadro's 
Hypothesis, 1 molecule of hydrochloric acid gas containi 
\ moL of hydrogen, 



.'. the hydrochloric acid molecule is HCl^. x is found 
from (ii) the vapour density. 

(iv) Electrolysis of concentrated hydrochloric^ acid solu- 
tion, using a carbon anode, yields hydrogen (cathode) and 
chloride (anode). After the liquid has become saturated with 
chlorine, it will be found that the volume of hydrogen which 
comes off from the cathode in a given time is exactly equal to 
the volume of chlorine evolved from 
the anode in that time. The same 
sort of argument as used in the pre- 
vious examples may be applied here. 
The composition by weight of 
hydrochloric acid is of great im- 
portance, since it yields the equiva- 
lent weight of chlorine (and also its 
atomic weight, since its valency is 1), 
upon which the atomic weights of 
so many elements depend. It has 
been determined by burning a known 
weight of pure hydrogen in pure 
chlorine and finding the weight of 
hydrochloric acid gas formed. By 
weighing also the chlorine used, the 
accuracy of the experiment could be 
checked. The student is advised to 
read the original papers in which 
this work is described. [Phil. Trans., 
205 A, 169, 1905 ; ibid., 209 A, 1, 

Oxides of Chlorine. Although the two elements do not 
combine together directly, three compounds of oxygen and 
chlorine are known : chlorine monoxide, C1 2 ; chlorine 
peroxide, C10 2 ; and chlorine heptoxide, C1 2 7 . They are 
all explosive and endothermic. 

Chlorine monoxide, C1 2 O. This substance was discovered 
in 1834 by BALAKD and is still best made by the method he 
employed, viz. the action of dry chlorine on dry mercuric 

FIQ. 111. 



oxide in the cold. The mercuric oxide is prepared by adding 
caustic soda to mercuric chloride solution 

HgCl a + 2NaOH 

= 2NaCl + H 2 + HgO ; 

it is washed and dried and a 
stream of cold dry chlorine is 
passed over it, when chlorine 
monoxide and mercuric oxy- 
chloride are formed 

2HgO + 2C1 2 

= HgO.HgCl a + Cl 2 0, 

Chlorine monoxide is an 
orange-yellow gas that con- 
denses to a reddish brown 
liquid if passed through a freez- 
ing-mixture. It explodes on 
heating and is a powerful 
oxidizing agent. It readily 
dissolves in water, most of it 
being converted into the weak 
hypochlorous acid, HC10, of 
which it is therefore the 

H 2 + C1 2 ^z 2HC10 

^=z 2H* + 2(310'. 

The composition of the gas 
was shown by BALABD and 
GAY-LUSSAC, by exploding a 
known volume and measuring 
the volumes of chlorine and 
oxygen produced. They found ^ 

that 2 volumes of chlorine ,, A _ 

. -, 01 t Fia - 112 - Electrolysis of 

monoxide gave 3 volumes of Hydrochloric Acid. 

mixed chlorine and oxygen; 

on treatment with caustic potash solution the chlorine 

dissolved and 1 volume of oxygen was left ; the volume of 


the chlorine was therefore 3 1=2 volumes. Therefore 
by Avogadro's Hypothesis 

2 molecules of chlorine monoxide give 2 t molecules of 
chlorine and 1 of oxygen. 

.'. 1 molecule of chlorine monoxide gives 1 molecule of 
chlorine and \ of oxygen. 

/. formula of the gas is C1 2 0. 

Chlorine peroxide or chlorine dioxide, C10 2 , was first 
prepared by DAVY in 1815, by the action of concentrated 
sulphuric acid upon potassium chlorate 

(i) KC10 8 + H 2 S0 4 = KHS0 4 + HC10,. 
(ii) 3HC10 8 = HC10 4 + 2C10 2 + H 2 0. 

, The action takes place in two stages, as shown above. First 
(i) the acid and chlorate react to give potassium hydrogen 
sulphate and chloric acid (HC10 3 ) ; this then decomposes (ii) 
into perchloric acid (HC10 4 ), water, and chlorine peroxide. 

BRAY'S method of preparing chlorine dioxide is to warm to 
60 a mixture of powdered potassium chlorate, oxalic acid, 
and a little water. It has the advantage of not being dan- 

Chlorine peroxide is a bright yellow gas with a character- 
istic smell. It is readily decomposed into its elements and is 
violently explosive. It can easily be liquefied, as was first 
shown by FARADAY in 1823. The liquid boils at 9 to 10 and 
freezes at 79 to orange-red crystals. 

Chlorine peroxide is a strong oxidizing agent ; phosphorus, 
sulphur, and many organic compounds such as sugar ignite 
spontaneously in the gas. The ignition of a mixture of 
potassium chlorate and sugar on addition of a drop of sulphuric 
acid is explained by the spontaneous combustion of a little of 
the sugar in the chlorine peroxide formed by the action of the 
acid on some of the chlorate ; ignition thus started rapidly 
spreads through the whole mixture. 

If some small pieces of phosphorus are placod under water 
and a few crystals of potassium chlorate added, the phosphorus 
may be made to burn under the water by introducing a little 
concentrated sulphuric acid through a thistle funnel on to the 


crystals of the chlorate ; chlorine peroxide is evolved and the 
phosphorus burns in it. 

Chlorine peroxide dissolves in water, yielding a solution 
which is acid, since chlorine peroxide combines with water to 
form a mixture of chlorous acid (HC10 2 ) and chloric acid 
(HC10 8 ) 

2C10 2 + H 2 O ^= HC10 8 + HC10 2 . 

It is therefore a mixed anhydride, like nitrogen peroxide (p. 
366). The aqueous solution of chlorine peroxide is unstable 
in the light, gradually decomposing into chloric acid, chlorine, 
and oxygen 

6C10 2 + 2H 2 = 4HC10, + Cl, + 2 . 

Chlorine peroxide will dissolve in solutions of caustic alkalis, 
forming chlorites and chlorates 

2C10 2 + 2KOH = KC10 2 + KC10, + H 2 0, 

aa would be expected. 

The constitution of chlorine peroxide was shown by DAVY 
and by GAY-LussAC by decomposing a known volume of the 
gas into its elements and measuring the volumes of the 
chlorine and oxygen produced. They found that 2 volumes 
of chlorine peroxide gave 1 volume of chlorine and 2 of 
oxygen, whence, by Avogadro's Hypothesis, the formula for 
the gas must be C10 2 . 

Chlorine heptoxide, C1 2 7 . This is the anhydride of 
perchloric acid, HC1O 4 . It was first made in 1900 by MICHAEL 
and CONN, by addition of anhydrous perchloric acid to phos- 
phorus pentoxide kept at a temperature of 10 by means 
of a freezing-mixture. 

2HC10 4 + P 2 O 6 = 2HP0 8 + C1 2 7 - 

After standing for 24 hours, at the same temperature, the 
mixture is heated to 82 when the chlorine heptoxide distils 
over as a colourless volatile oil. 

Chlorine heptoxide is extremely explosive, and its pre- 
paration involves considerable danger. It is unstable, and 
gradually decomposes on standing. It is not so powerful an 
oxidizing agent as the other chlorine oxides, e.g. it does not 


ignite sugar. With water it slowly reacts, to form perchloric 

C1 2 7 + H 2 = 2HC10 4 . * 

*' Euchlorine " is a mixture of chlorine peroxide and 
chlorine, made by the action of hydrochloric acid upon potas- 
sium chlorate. The compbsition of the mixture is variable. 
It was first prepared by DAVY (1811), who inclined to the view 
that it was a definite compound, but PEBAL (1875) proved 
that it was a mixture. 
Oxyacids of chlorine. 

Chlorine forms four oxy acids ; anhydrides are known of 
only two of them. 

Acid. Acid Anhydride. 

Hypochlorous acid, HC1O. Chlorine monoxide, C1 2 O. 

i -i TT^I^ TT i (CIO a i s the 

Chlorous acid, HC10,. Unknown) : , , 

_.. _ . . , Vr^n^ TT i "i mixed anny- 

Chlonc acid, HC10 8 . Unknown^ dride J 

Perchloric acid, HC10 4 . Chlorine hept oxide, C1 2 O 7 . 

Hypochlorous acid, HC10. The pure acid has not been 
isolated ; an aqueous solution of it is made by dissolving 
chlorine monoxide in water 

C1 2 O + H 2 ;z= 2HC10. 

On attempting to concentrate the solution, decomposition 
occurs, the hypochlorous acid splitting up into oxygen and 
hydrochloric acid 

2HC1O = 2HC1 + 2 . 

Hypochlorous acid in aqueous solution may also be made 
by treating bleaching- powder cautiously with the calculated 
quantity of dilute nitric acid and distilling under reduced 

The nearest approach to pure hypochlorous acid yet 
recorded is that of GOLDSCHMIDT., who prepared a 25 per 
cent, solution by the distillation under reduced pressure of a 
mixture of chlorine hydrate and yellow mercuric oxide. 

An aqueous solution of hypochlorous acid is yellow in colour 
and has the smell of chlorine monoxide. It is unstable, 
readily giving off oxygen, according to the equation given 


above. Hypochlorous acid is a weak, monobasic acid, but is 
a powerful oxidizing agent owing to the readiness with which 
it loses oxygen. It probably has the constitution represented 
by the formula H Cl. 

The salts of hypochlorous acid, the hypochlorites, are very 
important since they are largely used in the bleaching indus- 
try. The solution of potassium chloride and hypochlorite 
made by passing chlorine into a cold dilute solution of caustic 

2KOH + C1 2 = KC1 + KC10 + H 2 0, 

is called eau de Javelle, and has been in use as a bleaching 
agent since 1789. The corresponding solution made from 
caustic soda was first used in 1820 and was called eau de 
Labarraque. The bleaching action of hypochlorites in the 
presence of dilute acids is an oxidation process, as mentioned 
previously (p. 504). 

Like the acid itself, hypochlorites slowly decompose in 
aqueous solution, with liberation of oxygen ; the decompo- 
sition is hastened by the addition of cobalt or nickel salts as 

2NaC10 = 2NaCl + 2 , 
or, to show the reaction more generally 
2C10' = 2Cr + O 2 . 

An important derivative of hypochlorous acid is bleaching 

powder, Ca<^ . This is described in the section on 

X C1 

calcium (p. 248). It can be regarded as a " mixed salt," half 
a chloride and half a hypochlorite, as shown by the formula. 

Hypochlorites are difficult to prepare in the solid state, 
as they are so unstable ; however sodium hypochlorite > 
NaOC1.6H 2 0, was prepared in 1898 by MUSPRATT and 
SMITH, and calcium hypochlorite, Ca(OCl) 2 . 4H 2 O, by KINGZETT 
in 1875. 

They are white or slightly yellow crystalline solids. 

Chlorous acid, HC10,. This acid, like hypochlorous acid, 


is known only in solution. Mixed with chloric acid, it is 
formed when chlorine peroxide is dissolved in water 
2C10 2 + H 2 - HC10, + HC10 2 . 

The chlorites of certain metals have been isolated ; they are 
generally colourless, deliquescent, crystalline solids readily 
soluble in water. Sodium and potassium chlorites will bleach 
vegetable colours. The silver and lead salts are yellow, and 
only sparingly soluble. They explode when heated. A mix- 
ture of chlorite and chlorate is formed when chlorine peroxide 
is dissolved in a solution of caustic alkali (p. 513). 

Chloric acid, HC10 3 , can be obtained in aqueous solution 
by addition of the calculated weight of sulphuric acid to a 
solution of barium chlorate, and removing the precipitated 
barium sulphate 

Ba(C10 8 ) 2 + H 2 S0 4 = BaSO< + 2HC10 3 . 

By careful evaporation of the solution under reduced 
pressure it may be concentrated to a colourless syrupy liquid 
containing 30 to 40 per cent, chloric acid. The pure acid has 
not been prepared, since on further evaporation of the solution 
decomposition occurs, perchloric acid, oxygen and chlorine 
being formed 

8HC10 8 = 4HC10 4 + 30 2 + 2Cl a + 2H 2 0. 

The aqueous solution of the acid is a powerful oxidizing 
agent ; it sets fire to paper and many other organic sub- 
stances, and converts phosphorus into phosphoric acid, 
sulphur into sulphuric acid, etc. The constitution of the acid 

is still unsettled ; it may be HO Cl< I , since other com- 


pounds are known in which chlorine is tervalent. 

The salts of chloric acid, the chlorates, are stable com- 
pounds, some of them being of considerable importance both 
in the laboratory and in industry. 

Potassium chlorate, KC10 8 , is described in the chapter on 
potassium, p. 213. 

Barium chlorate, Ba(C10 8 ) 2 , and calcium chlorate, Ca(C10 t ) 2 , 


are white crystalline solids made by passing chlorine into a 
hot paste of the metallic hydroxide 

6Ca(OH) 2 + 6C1 2 = 5CaCl 2 + Ca(C10 3 ) 2 + 6H 2 0. 

The chlorate may be separated from the chloride by frac- 
tional crystallization ; the chlorate, being much the less 
soluble salt, separates out first. 

Chlorates decompose on heating, into the corresponding 
chlorides and oxygen 

Ba(C10 8 ) 2 = Bad 2 + 3O a . 
2KC1O 3 = 2KC1 + 30 2 . 

They are good oxidizing agents, and are used commercially 
in the match industry, for making fireworks, and in the 
manufacture of dyes. They are manufactured chiefly by 
electrical processes (p. 213). 

Perchloric acid, HC10 4 . This is the only oxyacid of 
chlorine which has been prepared in the pure state. It is 
made by distilling under reduced pressure a mixture of 
potassium perchlorate and concentrated sulphuric acid 
KC10 4 + H 2 S0 4 = KHS0 4 + HC1O 4 . 

The distillate of perchloric acid is purified by redistillation 
at a pressure of about 50-60 mm. 

Perchloric acid is a colourless oily liquid that fumes 
strongly in moist air and is very hygroscopic. It readily 
dissolves in water, with evolution of heat ; the solution is 
stable, but the anhydrous acid explodes if heated sufficiently 
or if dropped upon organic material such as wood or charcoal. 
It is an oxidizing agent, but not so active as chloric acid in 
this respect. 

The perchlorates are well known and are quite stable. They 
may be made by neutralizing the acid with the hydroxides or 
oxides of the metals, or by double decomposition of barium 
perchlorate with a metallic sulphate, etc., and also electro- 
lytically. Potassium perchlorate, KC10 4 , is formed by careful 
heating of potassium chlorate. The chlorate is heated until 
it melts ; shortly afterwards oxygen begins to come off, but 
after a time the evolution of gas ceases and the liquid becomes 
viscous. At this stage the fused mass consists of a mixture 


of potassium perchlorate and chloride, which may be separated 
from one another by fractional crystallization 

4KC10 3 - 3KC10 4 + KC1 * 

(or, since oxygen is evolved, 2KC10 3 KC10 4 + KC1 + 2 ). 
Perchloric anhydride is chlorine heptoxide, p. 513. 


Group in Periodic System : VII ; Symbol : Br ; Valency : 
I ; Atomicity of Vapour : 2 ; Atomic Weight : 79-92 ; Boiling 
Point : 59 ; Specific Gravity : 3-19. 

History. Bromine was .first made in a salt factory in 
Germany, about the year 1820, and sent to the great German 
chemist LIEBIG for examination ; he thought, however, that 
it was iodine chloride, and did not examine it very closely. 
The honour of the discovery of bromine must therefore be 
given to BALARD, who obtained it in 1826 from the mother- 
liquors (" bittern ") left after the recrystallization of salt 
from sea- water. He passed chlorine through these mother- 
liquors (which contain the bromides of sodium, potassium and 
magnesium) and noticed that a yellowish brown colouration 
was produced. Further experiments enabled him to isolate 
the coloured substance, which he obtained in the form of a dark 
brown liquid with a pungent smell. Balard, like Liebig, at first 
regarded it as iodine chloride, but failing to detect any iodine 
in it, he assumed it to be an element and called it originally 
muride but afterwards bromine (Greek, /fgo^uog, a stench). 

Occurrence. Bromine does not occur free in nature. It 
is found in the form of bromides, chiefly those of sodium, 
potassium, and magnesium. These are contained in sea- 
water and in certain mineral springs, and occur also in the 
solid state in the Stassfurt deposits in North Germany. 

Preparation. Bromine may be prepared in the labora- 
tory by heating a mixture of potassium bromide, manganese 
dioxide and sulphuric acid in a retort ; bromine distils over 
and is collected in a cooled receiver 

2KBr+Mn0 2 +3H 2 S0 4 =2KHS0 4 +MnS0 4 +2H 2 0+Br f . 


Other bromides may be used, but the potassium salt is the 
one usually met with. 

On the commercial scale, bromine is prepared by the action 
of chlorine upon the mother-liquors obtained in the recry- 
stallization of the salts of the Stassfurt deposits. Since the 
war cut off the supplies of bromine from Germany, the manu- 
facture of bromine in America (Ohio and the Saginaw Valley, 
Michigan) received a great impetus, and much bromine is now 
obtained commercially from the mother-liquors from the 
American springs. Bromine has also recently been extracted 
from sea-water on a commercially profitable scale. 

The Stassfurt and Ohio mother-liquors contain the bromides 
of potassium and magnesium. The hot mother-liquors are 

FIG. 113. Preparation of Bromine. 

allowed to flow down a tower filled with earthenware balls ; 
here they meet an upward stream of chlorine, and bromine 
is liberated 2KBr + C1 2 = 2KC1 + Br 2 . 

The bromine vapour and excess of chlorine escape from 
the top of the tower and are passed through a spiral cooling- 
tube surrounded by water, where most of the bromine con- 
denses and may be collected in a receiver. Any uncondensed 
bromine is trapped by passing through a small tower contain- 
ing wet iron filings, where iron bromide is formed ; this is 
used for making potassium bromide (p. 525). The spent 
mother-liquors that leave the bottom of the main tower 
carry a little of the chlorine and bromine with them ; they 
are therefore allowed to circulate over sandstone shelves in a 
small chamber through which steam is blown. The chlorine 
and bromine are thus blown out of the solution and, mixing 



with the main stream of chlorine from the chlorine generator, 
pass up the tower again. 

The chief impurities in the bromine obtained in the com- 
mercial process are chlorine and water. The water may be 
removed by shaking with concentrated sulphuric acid ; if the 
dry bromine is then distilled over potassium bromide or iron 
bromide it is freed from chlorine, which replaces the bromine 

Hot mother liquors 


FIQ. 114. Manufacture of Bromine. 

in the bromide added and is left behind as potassium or iron 

2KBr + C1 2 - 2KC1 + Br 2 . 

Iodine, which sometimes occurs as an impurity, is removed 
by addition of a little starch paste, which absorbs it (p. 531). 

The purest bromine hitherto obtained was made by SCOTT 
in 1913. [See SCOTT, Trans. Chem. Soc., 131, 847, 1913.] 


Properties. Bromine is a reddish brown, heavy, volatile 
liquid, which boils at 59, freezes at 7-3, and has a specific 
gravity of 3*19. It readily passes into vapour at ordinary 
temperatures ; the vapour is red in colour and very irritating 
to the mucous membrane and the eyes. Liquid bromine 
produces yellow burns when dropped on the skin. The 
vapour density of bromine vapour shows that at first the 
molecules of the gas are Br 2 (V.D. = 79-92) ; as the temper- 
ature rises, however, partial dissociation occurs. 
Br 2 ; * Br + Br. 

Bromine dissolves slightly in water ; the solution is called 
bromine-water, and on cooling to deposits red crystals of 
bromine hydrate, Br 2 .8H 2 O. (Compare chlorine hydrate,, 
p. 504.) Bromine is much more soluble in organic solvents 
such as chloroform, ether, carbon disulphide and benzene. 
It is absorbed by the siliceous earth called kieselguhr ; the 
product gradually gives off bromine on exposure to air and is 
sold as a disinfectant under the (incorrect) name of bromum 
solidificatum, or solidified bromine. 

In its chemical properties bromine closely resembles 
chlorine ; it combines directly with hydrogen (to form hydro- 
bromic acid gas), with metals (to form salts, the bromides 
" halogen "), and with phosphorus (to form phosphorus tri- 
and penta- bromides). The reaction between metallic potas- 
sium and liquid bromine is explosive. 

Owing to its affinity for hydrogen, bromine is an oxidizing 
agent, though less powerful than chlorine. It will bleach, 
though not very energetically, and turns starch -solution yellow 
(distinction from iodine, p. 531). Ordinary bromine consists 
of 2 isotopes, atomic weights 79 and 81. 

The atomic weight of ordinary bromine is 79-92 ; it was 
determined by STAS by converting a known weight of pure 
silver into silver bromide, which was weighed, and by GUYB 
and also by MOLES, from the vapour density of pure hydro- 
bromic acid gas. 

Bromine is required in large quantities (several million 
pounds per annum) for addition in the form of ethylene 
dibromide, C 2 H 4 Br 8 to " ethyl petrol." About 2 c.c. 



ethylene dibromide is added to every gallon of petrol, its 
function being to provide bromine to convert the lead of the 
lead tetraethyl (p. 330) into lead bromide. The lead would 
otherwise foul the sparking-plugs. 


Hydrogen bromide, or Hydrobromic acid,HBr. When 
hydrogen and bromine vapour are heated together in equal 
volumes, combination takes place with measurable velocity 
at about 250, and at higher temperatures proceeds much 
more quickly and practically completely. At still higher 
temperatures, however, the hydrogen bromide dissociates 

H 2 + Br 2 ^=i 2HBr. 

The best yield of hydrobromic acid, by the synthetic 
method, is therefore obtained by passing the mixture of 

FIG. 115. Preparation of Hydrobromic Acid. 

hydrogen and bromine vapour over a catalyst heated to a 
moderately high temperature ; e.g. by passing the mixture 
through a tube containing a platinum spiral electrically 

Hydrobromic acid cannot be obtained pure by heating 
sulphuric acid with a bromide (cf. preparation of hydro- 
chloric acid), since it is a mild reducing agent, and on 
liberation reduces the hot sulphuric acid to sulphur dioxide, 
being itself oxidized to bromine 

i. KBr + H 2 S0 4 = KHS0 4 + HBr. 
ii. H 2 S0 4 + 2HBr = 2H 2 O + S0 2 + Br 2 , 

or 2KBr + 3H 2 S0 4 = 2KHSO 4 + 2H 2 + S0 2 + Br a . 
Some of the hydrobromic acid, it is true, escapes decompo- 
sition, but the mixture of gases evolved, hydrobromic acid, 


sulphur dioxide and bromine, cannot be conveniently 

Hydrobromic acid is therefore generally prepared by the 
action of water on phosphorus tribromide, PBr 3 

PBr, + 3H 2 = H 3 P0 3 + 3HBr. 


Similar reactions with the trichloride and tri-iodide of 

FIG. 116. Preparation of Hydrobromic Acid. 

phosphorus give hydrochloric acid and hydriodic acid ; 
hydriodie acid indeed is prepared in this way, but the method 
is not used for hydrochloric acid since simpler ones are 

Red phosphorus, water, and sand arc stirred up together 
into a thin paste and put in a flask fitted with a cork carrying 
a dropping-funnel and deli very- tube. The deli very -tube IB 
connected to a U-tube containing bits of broken glass smeared 
with a paste of red phosphorus and water. Bromine is placed 
in the funnel and allowed to run in slowly. As each drop 


touches the phosphorus a bright flash occurs, and hydro- 
bromic acid gas is given off. The bromine combines with the 
phosphorus to form phosphorus tribromide, and this is then 
decomposed by the water present, according to the equation 
given above. 

Owing to the volatile nature of bromine, the gas that 
comes off from the delivery tube carries with it some bromine 
vapour ; this is removed in the U-tube, where it is converted 
into hydrobromic acid by the water and phosphorus. The 
U-tube is, in fact, a secondary apparatus for producing hydro- 
bromic acid. 

The hydrobromic acid gas is then collected by downward 
displacement or over mercury, or if a solution of the hydro- 
bromic acid is required, as in Fig. 109. 

If larger quantities of the gas are required, it is better to 
use ERDMANN'S method, which is very simple to carry out. 
It consists in the addition of bromine to benzene (C 6 H 6 ) in 
the presence of a catalyst, such as reduced iron or aluminium 
powder, when the following reaction occurs 

C 6 H, + 2Br 2 = C 6 H 4 Br a + 2HBr. 


The benzene (which should be dry) is placed in a flask and 
a little reduced iron is added. Bromine is then gradually 
run in, when the action starts at once ; it may be necessary 
to cool the flask, as the reaction is vigorous. The hydro- 
bromic acid gas that is evolved carries over benzene vapour 
and bromine vapour with it ; it is freed from the latter by 
passing through a U-tube containing ferric bromide, FeBr 8 , 
and from the former by anthracene contained in a second 

The aqueous solution of hydrobromic acid may be made by 
dissolving the gas in water (taking the precautions usual in 
the case of a very soluble gas), or directly in one of the follow- 
ing ways 

(i) By passing sulphuretted hydrogen through bromine- 
water, and filtering off the precipitated sulphur 

H 2 S + Br 2 = 2HBr + S. 


(ii) By passing sulphur dioxide through bromine-water, or 
bromine covered by a layer of bromine-water 

S0 2 + Br 2 + 2H 2 = H 2 S0 4 + 2HBr. 

From the mixture of dilute sulphuric and hydrobromic 
acids obtained in this way, dilute hydrobromic acid may be 
made by distillation. 

Properties. Hydrobromic acid is a heavy colourless acid 
gas that fumes in moist air, dissolves readily in water giving 
a strongly acid solution, and in its general properties closelj 
resembles hydrochloric acid. It can be condensed to a 
colourless liquid boiling at 69 and freezing to a colourless 
crystalline solid at 86. 

The aqueous solution is highly ionized 

HBr ;r IT + Br'. 

Hydrobromic acid, like hydrochloric acid (p. 507), forms a 
constant-boiling mixture with water. This at atmospheric 
pressure boils at 126 and contains 47-8 per cent. HBr by 
weight. In aqueous solution hydrobromic acid slowly decom- 
poses on exposure to air and light, the solution turning yellow 
owing to liberation of bromine by oxidation 
O a + 4HBr = 2H 2 + 2Br 2 . 

A convenient way of obtaining the gas from the aqueous 
solution is to add an excess of anhydrous calcium bromide, 
or to drop the solution from a dropping funnel on to the 
anhydrous salt. 

The salts of hydrobromic acid are called bromides. The 
most important is potassium bromide, KBr, which finds many 
applications in the laboratory, in medicine, and in photo- 
graphy. It is made from the iron bromide obtained in the 
manufacture of bromine (p. 519) by dissolving it in water and 
adding potassium carbonate solution ; an iron hydroxide 
is precipitated, carbon dioxide is evolved, and a solution of 
potassium bromide is left from which the*salt may be crystal- 
lized on evaporation. 

2FeBr 8 + 3K 2 C0 3 + 3H 2 = 2Fe(OH) 8 + 6KBr + 3CO t , 
As ths iron bromide from the bromine works does not yield 


sufficient potassium bromide to satisfy the demand, more is 
made by adding bromine to iron filings and is then converted 
into potassium bromide by the method described. 

Potassium bromide may also be made by dissolving bromine 
in concentrated caustic potash solution, evaporating to dry- 
ness, and heating the residual mass of bromide and bromate 
with charcoal, or even alone 

3Br 2 + 6KOH - 5KBr + KBrO 3 + 3H 2 0. 
5KBr + KBrO, = GKBr + oxygen, 
or, with charcoal, 

5KBr + KBr0 3 + 3C = GKBr + SCO. 

It is a white crystalline solid, soluble in water and is used 
in medicine as a sedative or sleeping-draught. 

All metallic bromides are soluble in water except those of 
silver, lead and mercury (ous). The formation of pale-yellow 
insoluble silver bromide is used as a test for a bromide in 
qualitative analysis. Silver bromide is used in photography 
(p. 233). 

Oxides of Bromine are unknown. 

Oxyacids of Bromine. Hypobromous acid, HBrO, and 
bromic acid, HBr0 3 , are known, and there is some evidence 
for the existence of bromous acid, HBr0 2 . 

Hypobromous acid, HBrO. BALABD prepared an aqueous 
solution of hypobromous acid by shaking bromine-water with 
yellow mercuric oxide. By further addition of bromine and 
mercuric oxide a solution can be obtained containing 6 to 
7 per cent, of hypobromous acid. Distillation of this liquid 
under reduced pressure yields a pale yellow dilute solution of 
the acid. 

2Br 2 + 2HgO + H 2 = 2HBrO + HgO . HgBr 2 . 

The pure acid has not been isolated ; attempts at further 
concentration of the solution cause the acid to decompose into 
water, bromine, and oxygen. 

Hypobromites, mixed with bromides, are formed when 
bromine is dissolved in cold dilute solutions of caustic alkalis 

2KOH + Br 2 = KBr + KBrO + H E 0. 


They are very similar to the hypochlorites, but are even 
less stable ; they have bleaching properties, and will liberate 
nitrogen from urea, for the estimation of which they are of ton 

Bromous acid, HBr0 2 . According to RICHARDS (1906) 
a solution of bromous acid is formed when a large excess of 
bromine water is added to silver nitrate solution. 

Bromic acid, HBr0 3 . This acid is known only in dilute 
solution and in the form of its salts, the bromates. The 
aqueous solution of the acid has been prepared in many ways, 
e.g. by the action of bromine on a suspension of silver bromate 
in hot water (KAMMERER, 1869) 

5AgBr0 8 + 3Br 2 + 3H 2 O = 5AgBr^ + 6HBr0 3 , 

or by addition of dilute sulphuric acid to a solution of barium 
bromate (BALARD, 1826) 

Ba(Br0 3 ) 2 + H 2 SO 4 = BaS0 4 + + 2HBrO,. 

The aqueous solution can be concentrated by distillation 
under reduced pressure to a syrupy colourless liquid contain- 
ing 50-6 per cent. HBrO,. Further concentration brings about 

4HBr0 3 = 2H 2 + 50 2 + 2Br 2 . 

Bromic acid is a strong acid and a powerful oxidizing agent ; 
thus it oxidizes hydrobromic acid to bromine 

HBr0 3 + 5HBr = 3H 2 + Br,, 
nitrous acid to nitric acid 

5HN0 2 + 2HBr0 3 - 5HN0 3 + H 2 + Br 2 , 
and sulphur dioxide solution to sulphuric acid 

2HBr0 3 + 5H 2 S0 3 = Br 2 + H 2 + 5H 2 S0 4 . 

The bromates show close similarity to the chlorates. A 
mixture of potassium bromide and bromate is made by adding 
bromine to hot concentrated caustic potash solution 

3Br 2 + 6KOH = 5KBr + KBr0 8 + 3H 2 0. 
The silver salt, being insoluble, may be prepared from a 


solution of the potassium salt by precipitation. Bromatea 
yield bromides and oxygen on heating. 

Perbromic acid and perbromates do not exist. 


Oroup in Periodic System : VII ; Symbol : I ; Valency : 1 j 
Atomicity of Vapour : 2 ; Atomic Weight : 126-93 ; Melting 
Point: 116-1; Specific Gravity : 4-93. 

History. In 1812 a French saltpetre manufacturer, 
COURTOIS, found that on adding concentrated sulphuric acid 
to the mother-liquor of the sodium carbonate extracted from 
the ashes of seaweed, a black powder was precipitated, which 
on heating was converted into a violet vapour. Courtois sent 
some of his new substance to GAY-LUSSAO, who made a 
thorough investigation of it and showed that it was an element- 
similar in many respects to chlorine. He called it iodine, 
from the violet colour of its vapour (Greek, IQSI&YIS), and 
succeeded in preparing hydriodic acid, HI, from it. SIR 
HUMPHRY DAVY, who was passing through Paris at the 
time, also made investigations on the new element, a specimen 
of which was given to him by AMPERE, and confirmed Gay- 
Lussac's results. 

[For an interesting account of the history of iodine, see 
Chemical News, 99, 193 (1909).] 

Occurrence. Iodine is not found free in nature. It is, 
however, widely distributed (although in small quantities) in 
the form of iodides, and to a less extent, iodates. It occurs 
as sodium iodate, NaIO 3 , mixed with sodium nitrate, in Chile 
saltpetre or caliche (about 0-3 per cent. NaIO 3 }, and as the 
iodides of sodium, magnesium and potassium in sea-water. 
It is an interesting fact that seaweeds have the power of 
absorbing iodine from sea-water ; thus while the percentage 
of iodine in the sea is 0-001, the percentage in the dry matter 
of Laminaria, a deep-sea seaweed, is nearly 0*5. Sponges 
seem to possess this power to an astonishing degree ; accord- 
ing to HUNDESHAGEN, tropical sponges may contain as much 
as 14 per cent, of iodine ! 



An organic compound of iodine is a constant constituent of 
the important thyroid gland in man (p. 531). 

Extraction. Iodine is obtained both from seaweeds and 
from caliche. 

(i) Extraction from Seaweed. The seaweed is dried during 
the summer months and then burnt, forming an ash known 
as kelp in this country and varech in Normandy. The kelp 
contains the sulphates, carbonates, chlorides, bromides and 
iodides of sodium and potassium, mixed with carbon. It is 
stirred up with hot water in iron pots, and the insoluble 
matter allowed to settle. The solution is then run off and 


FIG. 117. Manufacture of Iodine from Kelp. 
A. Iron Pot. BB. Udells. 

concentrated, when the sulphates, carbonates and chlorides 
separate out, being less soluble, while the bromides and iodides 
are left in the mother-liquor. (The proportion of bromides to 
iodides is small.) 

The mother-liquor is mixed with manganese dioxide and 
concentrated sulphuric acid and heated in cast-iron pots, 
when the iodine distils over and is collected in specially- 
shaped condensers called alitdels or udells. Towards the end 
of the reaction, bromine conies over, from the bromides 
present. It is collected, if present in sufficient quantity, 
otherwise it is allowed to go to waste. 
* The iodine obtained in this way is washed, dried, and 
purified by sublimation over potassium iodide. 

(ii) Extraction from Caliche. The crude sodium nitrate is 


recrystallized from water, when the sodium iodate, NaIO 8 , 
is left in the aqua vieja or mother-liquor. This is run into 
wooden or lead-lined vats and mixed with the calculated 
weight of sodium bisulphite, NaHS0 3 , to precipitate all the 
iodine, according to the equations 

(i) NaHS0 3 + NaI0 3 - Na 2 S0 3 + HI0 3 . 
Here the acid sodium sulphite liberates iodic acid from some 
of the iodate. 

(ii) 3NaHS0 3 + HI0 3 = 3NaHS0 4 + HI. 
The reducing properties of the sulphite convert iodic acid to 
hydriodic acid. 

(in) 5HI + HI0 3 = 3H 2 + 3I 2 ^ . 

Iodic and hydriodic acids cannot exist together in solution. 
They immediately react to form iodine and water. 

The black precipitate of iodine is filtered off, and purified as 
above, while the remaining liquid is used over again for the 
recrystaUization of a fresh lot of caliche. 

Modifications of this process are also employed in certain 
districts in South America. 

In the laboratory iodine is generally made by heating a 
mixture of manganese dioxide and sulphuric acid with potas- 
sium iodide, in a retort. Iodine sublimes over into the neck 
of the retort. 
2KI + Mn0 2 + 3H 2 S0 4 = 2KHS0 4 + MiiS0 4 + 2H 2 + I 2 . 

Properties. Iodine is a lustrous black crystalline solid, 
which sublimes if quickly heated, but if slowly heated melts 
at 116. The liquid boils at 184 to a deep purple vapour, 
whose density at the boiling-point corresponds to the formula 
I 2 . At 600 dissociation begins, and at 1,600 is complete, 
the molecules of the vapour then all being monatomic 

i, ^= i + 1. 

Iodine is only slightly soluble in water, but is much more 
soluble in organic solvents. It is also very soluble in potas- 
sium iodide solution ; a certain amount of combination 
occurs, potassium tri-iodide being formed 
KI + I 2 ^z KI,. 


This substance has been isolated in the form of black crystals. 
It loses a molecule of iodine so readily in aqueous solution 
that a solution of iodine in potassium iodide may be regarded 
in volumetric analysis as containing as much free, iodine as 
was dissolved. This fact is very convenient, since concen- 
trated aqueous solutions of free iodine cannot be prepared in 
the usual way owing to the small solubility of the element. 

Iodine has the usual halogen properties, but is less reactive 
than chlorine and bromine ; however, a mixture of yellow 
phosphorus and iodine takes fire spontaneously, forming 
iodides of phosphorus. Iodine will also combine directly 
with many metals, forming salts, the iodides. 

The most characteristic property of iodine is perhaps the 
reaction with starch. When iodine is added to starch solution 
a deep blue colour is produced ; this reaction is so sensitive 
that 0-00001 grams of iodine in a litre of water may be detected 
by it. On heating, the blue colour disappears, but it reappears 
on cooling. It is permanently destroyed by addition of 
ammonia. The composition of the blue substance is un- 
known ; it is probably an " adsorption compound " of iodine 
in starch, and is almost certainly not a true chemical com- 

Iodine is essential to the well-being of the human body. 
Its absence may lead to goitre or cretinism. Though normally 
supplied to the body in certain foods, e.g. spinach and lettuce 
among vegetables and butter and milk among animal foods, 
it is well to add small quantities of sodium iodide to table salt. 
Many salt manufacturers now make this addition as a matter 
of course. The body of a normal adult contains about 20 
milligrams of iodine, approximately half the total quantity 
being present as the compound ihyroxin, in the thyroid gland. 

Hydriodic acid, or Hydrogen iodide, HI. If a 

mixture of hydrogen and iodine vapour is heated, partial 
combination takes place and an equilibrium between hydrogen, 
iodine, and hydrogen iodide is set up 

H 2 -f I 2 ^= 


At 448, the mixture contains 75 per cent. HI, but the rate 
of formation at this temperature is 'so slow that it is best to 
add a catalyst, such as platinum black, to accelerate the 

Hydriodic acid gas, however, is generally prepared by the 
action of water upon phosphorus iodide 

PI 3 + 3H 2 = H 3 P0 3 + SHI. 

In practice, an apparatus similar to that employed for the 
preparation of hydrobromic acid (p. 523) is used. Water is 
placed in the dropping-funnel and a mixture of red phosphorus 
and iodine in the flask. On adding the water, drop by drop, 
hydriodic acid gas comes off. Since the actual formation of 
phosphorus iodide in this reaction has not been observed, it 
may be better to write the equation as 

2P + 3I 2 + 6H 2 = 6HI + 2H 3 P0 3 . 

As iodine is not very volatile at ordinary temperature, 
there is no necessity to have the U-tube on the delivery tube. 

The hydriodic acid gas that comes off is collected by 
downward displacement (not over mercury, which it attacks). 

An aqueous solution of hydriodic acid may be made by 
dissolving the gas in water (precautions for very soluble gas), 
or by passing sulphuretted hydrogen through iodine sus- 
pended in water 

I 2 + H 2 S - 2HI + S. 
(Cf. Br 2 + H 2 S = 2HBr + S, and C1 2 + H 2 S = 2HC1 + S.) 

Hydriodic acid cannot be made by the action of concentrated 
sulphuric acid upon potassium iodide, since it is a reducing 
agent, and reduces the acid to sulphur, sulphuretted hydrogen 
and sulphur dioxide, being itself oxidized to iodine. 

Properties. Hydrogen iodide is a colourless heavy acid gas, 
fuming in moist air, and readily soluble in water, giving a 
strongly acid solution. It can be condensed to a colourless 
liquid, boiling at 35-5 and freezing at 51 to a colourless 
crystalline solid. 

In general chemical properties it closely resembles hydro- 
chloric and hydrobromic acids. Unlike them, however, it is 


an endothermic substance and a powerful reducing agent ; 
hydrochloric acid and hydrobromic acid are both strongly 
exothermic, and are not reducing agents under ordinary 
circumstances, although hydrobromic acid may sometimes 
show reducing powers (p. 522). 

Like the other halogen acids, hydriodic acid forms a constant- 
boiling mixture, that formed at atmospheric pressure boiling 
at 127 and containing 57 per cent. HI. Aqueous hydriodio 
acid is used in organic chemistry as a reducing agent. When 
freshly prepared it is colourless, but rapidly goes brown 
owing to oxidation of the acid with liberation of iodine 

2HI + oxygen = H 2 + I 2 . 

The iodides are very similar to the chlorides and bromides. 
The formation of the insoluble yellow silver salt is used as a 
test for iodides in solution. 

Oxides of Iodine. 

Iodine forms three oxides iodine dioxide, 10 2 , iodine 
pentoxide, I 2 5 , and " iodine iodate," I 4 2 . 

Iodine dioxide, I0 2 (or I 2 4 ), was first prepared by 
MILLON (1844). It may be made by heating sulphuric acid 
with iodic acid, or by the action of concentrated nitric acid 
upon dry powdered iodine in the cold. It is a yellow crystal- 
line solid that may be purified by recrystallization from 
concentrated sulphuric acid. It reacts with water to form 
iodic acid (HI0 3 ) and iodine 

5I 2 O 4 + 4H 2 = 8HI0 8 + I,. 

Iodine dioxide, unlike chlorine dioxide, is not explosive. 

Iodine pentoxide, I 2 O 6 , is the anhydride of iodic acid, 
from which it may be made by the action of gentle heat 

It is a white crystalline solid that dissolves in water to 
form iodic acid. When exposed to a dull red heat it splits up 
into its elements. It is a strong oxidizing agent, converting 
carbon monoxide into the dioxide even at ordinary tempera- 


tares, while on heating it will convert sulphur dioxide into 
the trioxide 

SCO + I 2 6 = 5C0 2 + I 2 . * 
5S0 2 + I 2 5 5S0 3 + I 2 . 

Oxyacids of Iodine. Iodine forms three oxyacids 
Hypoiodous acid, HIO. 
lodic acid, HI0 3 . 

Periodic acid, HI0 4 (or H 5 I0 6 , see below). 
Hypoiodous acid, HIO. When iodine is added to dilute 
caustic potash solution it forms a mixture of potassium iodide 
and potassium hypoiodite 

2KOH + I 2 = KI + KIO + H 2 O. 

Hypoiodous acid, however, is such a weak acid that unless a 
large excess of alkali is present the potassium hypoiodite is 
hydrolysed into Tiypoiodous acid itself and caustic potash 

KIO + H 2 O ;=z HIO + KOH. 

A solution of hypoiodous acid may also be obtained by 
shaking a solution of iodine in alcohol with yellow mercuric, 
oxide. (Compare hypochlorous and hypobromous acids, 
pp. 514 and 526.) 

The free acid has not been isolated. The aqueous solution 
is yellow in colour and has oxidizing and bleaching properties. 

lodic acid, HI0 3 , is made by heating iodine with a large 
excess of fuming nitric acid. The reaction is very slow, but 
is greatly accelerated if a stream of oxygon is bubbled through 
the reacting mixture. 

lodic acid is a white deliquescent crystalline solid. It was 
discovered by GAY-LUSSAC, who made it by passing a stream 
of chlorine through a suspension of iodine in water 

I 2 + 5C1 2 + 6H 2 = 2HI0 3 + 10HC1. 

It is a strong oxidizing agent, and also a strong acid, i.e. 
in aqueous solution is is highly ionized. Cryoscopic experi- 
ments on concentrated solutions of the acid indicate that it 
may have the double formula (HI0 3 ) 2 or H 2 I 2 O (5 ; the exist- 


ence of acid iodates is additional evidence for the double 

lodic acid and hydriodic acid react together in solution, 
forming iodine and water (see p. 530) 

5HI + HI0 3 = 3I 2 + 3H 2 O. 

The chief iodate is that of potassium, which is made by 
dissolving iodine in a hot concentrated solution of caustic 

3I 2 + 6KOH = 5KI + KI0 3 + 3H 2 0. 

It is a white crystalline solid, soluble in water, and splitting 
up into potassium iodide and oxygen on heating. (Cf. 
potassium chlorate and bromate.) 

2KIO 8 = 2KI + 30 2 . 

Periodic acid. Several periodic acids are known ; the 
chief are meta-periodic acid, HI0 4 , and para-periodic acid, 
H 5 I0 6 . The barium salt of the para acid is formed when 
barium iodate is carefully heated 

5Ba(I0 3 ) 2 - Ba 5 (I0 6 ) 2 + 4I 2 + 90 2 . 

The potassium salt of the meta acid can be obtained by 
electrolysis of an alkaline solution of potassium iodate. 

Free para-periodic acid may be made from the barium salt 
by addition of sulphuric acid 

Ba 6 (I0 6 ) 2 + 5H 2 S0 4 = 5BaS0 4 + + 2H 6 I0 6 . 

Evaporation of the filtered solution produces crystals of the 
para-acid. On heating para-periodic acid to 100 at low 
pressures, it yields the met a -acid - 

H 5 I0 6 - HIO 4 + 2H 2 0. 


1. Show that fluorine, chlorine, bromine and iodine are justifiably 
grouped together. 

2. Describe the preparation and properties of (a) fluorine, (6) hydro- 
fluoric acid. 

3. Give an account of the history of chlorine. 

4. How is chlorine prepared in the laboratory T What are its 
properties ? 

5. Describe the manufacture of chlorine. 


6. What is the action of chlorine upon alkalis ? 

7. Write an account of the oxides and oxyacids of chlorine. 

8. How is bromine manufactured ? What are its properties ? 

9. Describe the preparation and properties of hydrobromic acid gas. 

10. How does iodine occur in nature ? How is it manufactured ? 

11. Compare and contrast the behaviour of sulphuric acid with 
(a) potassium chloride, (6) potassium bromide, and (c) potassium iodide. 

12. How is iodic acid made ? What are its properties ? 

13. Discuss the valency of the halogens. 


GROUP VII, Sub-group A 


Until 1925, manganese was an orphan element, occupying 
a position of splendid isolation in the Periodic System, and 
although it was forced into Group VII with the halogens it 
must have felt extremely uncomfortable among those fiery 
and excitable neighbours. However, in 1925, the existence 
of two new elements, masurium and rhenium, was inferred 
spectroscopically, and these metals would complete Group 
VII, sub-group A. There is still doubt as to the identity of 
masurium, but rhenium has now been isolated and its atomic 
weight (188*7) determined, while certain of its compounds 
have been carefully studied. Potassium per-rhenate, KRe0 4 , 
is already included in the manufacturer's lists. Rhenium 
occurs in certain tantalum and platinum minerals. 


Group in Periodic System : VII ; Symbol : Mn ; Valency : 
2, 3, 4, 7 ; Atomic Weight : 54-93 ; Melting Point : 1,260 ; 
Specific Gravity : 7-2-8-0. 

History. The ancient chemists confused pyrolusite, 
Mn0 2 , with magnetite, Fe 3 4 , but PLINY recognized it as 
a distinct substance and called it magnesia nigra. It was 
well known to the Arab chemists of the Middle Ages, who 
carefully distinguished between it and magnetite and magnesia 
alba (MgO). It has been used for centuries in glass manu- 
facture, to take out the green colour caused by the presence 



of iron. Pyrolusite was for long considered to be an ore of 
iron, but in 1774 SOHBELB showed that it contained a new 
metallic element, and shortly afterwards the % metal was 
isolated by GAHN who reduced the pyrolusite with carbon. 

Occurrence. Maganese is found in nature chiefly as the 
dioxide, Mn0 2 , which is called pyrolusite. This mineral is 
found in Germany, Russia, India, Brazil, North America, 
Spain, the Caucasus and other countries. Other manganese 
ores are manganese blende, MnS ; Hausmannite, Mn 3 O 4 ; and 
manganese spinel, (Mn, Mg)(Fe, Mn) 2 O 4 . 

Preparation. Manganese is generally made nowadays 
either by reducing the oxide with powdered aluminium 
3Mn0 2 + 4A1 = 2A1 2 3 + 3Mn, 

or by the reduction of the dioxide with carbon in the electric 

MOISSAN'S method gives a purer product. It consists of the 
electrolysis of a solution of manganous chloride, MnCl 2 , by 
means of a powerful electric current. The cathode is of 
mercury and a manganese-mercury amalgam containing about 
4 per cent, of manganese is obtained. This is carefully washed 
and dried and the mercury distilled off under reduced pressure. 

Properties. Pure manganese is a hard, brittle, greyish 
metal with a reddish tinge. The specific gravity is 7-39-8-0 
according to the method of preparation ; probably 7-39 is 
the most nearly correct value hitherto obtained. The metal 
melts at 1,260 and boils at 1,900. It readily oxidizes in the 
air and is therefore kept under rock-oil, like sodium. It 
decomposes water in the cold with evolution of hydrogen, 
and is very soluble in dilute acids, less so in the concentrated 
acids. It combines directly with oxygen, nitrogen, chlorine, 
carbon, silicon, boron and phosphorus, when heated. 

The atomic weight of the element has been determined by 
analysis of manganous bromide, MnBr 2 , and chloride, MnCl 2 . 
The value generally accepted is 54-93. 

Manganese has recently become of great importance in the 
manufacture of alloys such as ferro-manganese, manganese 
bronze, spiegeleisen, cupro-manganese, manganin, etc. These 


find various uses ; manganese bronze, for example, is used 
for the propellers of ships while manganin is used in rheostat? 
or resistance coils. 

It is said that rats fed on a diet free from manganese lose 
all love for their offspring. 

COMPOUNDS OF MANGANESE. Manganese forms a whole 
family of oxides 

Manganous oxide, MnO, basic. Forms manganous salts. 

Trimanganic tetroxide, Mn 3 4 . Resembles Fe 3 4 . 

Manganic oxide, Mn 2 3 . Slightly basic, forming manganic 

Manganese dioxide, Mn0 2 . Slightly acidic, forming man- 

Manganese trioxide, MnO 3 . Acidic, forming manganates. 

Manganese hep t oxide, Mn 2 O 7 . This also is acidic, forming 
permanganates . 

Manganese forms two series of salts, the manganous salts 
corresponding to the oxide MnO, and the manganic salta 
corresponding to the oxide Mn 2 O 3 . In the first series man- 
ganese is bivalent and in the second tervalent. The common 
salts of manganese are manganous salts. 

Manganous oxide, MnO, is obtained as a bright green 
powder by heating manganese dioxide or manganous car- 
bonate in a stream of hydrogen. It is soluble in dilute acids, 
forming manganous salts. 

Manganous hydroxide, Mn(OH) 2 , is formed as a white 
precipitate (rapidly going brown in the air, owing to oxida- 
tion) when caustic soda or potash solution is added to a 
solution of a manganous salt. 

Manganous chloride is obtained in the form of pink crys- 
tals containing 4 molecules of water of crystallization, 
MnCl 2 .4H 2 O, by dissolving the metal, oxide, hydroxide, or 
carbonate in dilute hydrochloric acid and evaporating the 

Manganous carbonate, MnC0 3 , is obtained as a dirty white 
precipitate when sodium carbonate is added to a solution of 
a manganous salt. It is unstable in moist air, slowly oxidizing 
to manganic hydroxide, Mn(OH) t , and other substances. 


Manganese dioxide, Mn0 2 , is a black powder. The 
native form, pyrolusite, contains on an average about 80 
per cent. Mn0 2 . When pyrolusite is required for the manu- 
facture of chlorine it is important from the buyer's point 
of view to know how much manganese dioxide it contains. 
This is determined by taking a known weight of the finely 
ground ore and boiling it with hydrochloric acid, until all 
black particles have disappeared. 

Mn0 2 + 4HC1 = MnCl a + 2H 2 O + Cl a - 

The chlorine which is evolved is passed through potassium 
iodide solution, where it liberates an equivalent weight of 
iodine. The weight of iodine set free is estimated by titration 
with standard sodium thiosulphate. 

Manganese dioxide is a good oxidizing agent. It is used 
in the manufacture of chlorine, for decolourizing glass, as a 
depolarizer in Leclanche cells and dry batteries, and in the 
preparation of potassium and sodium permanganates. 

It will dissolve in concentrated alkalis, forming salts called 
manganites, such as CaMnO 8 , calcium manganite. (See 
Weldon's Process, p. 501.) 

Manganates and Permanganates. When manganese 
dioxide is fused with caustic potash or caustic soda in the 
presence of air or, better, with addition of a strong oxidizing 
agent such as potassium chlorate or nitrate, a dark green 
mass is obtained. This consists of impure sodium or potassium 

2Mn0 2 + 4KOH + O 2 = 2K 2 MnO 4 + 2H 2 0. 

If the mass is dissolved in water a green solution is formed 
from which dark green crystals of the manganate, K 2 MnO 4 , 
may be obtained by careful evaporation under reduced pres- 
sure. Potassium manganate is isomorphous with potassium 
sulphate and potassium chromate. 

On warming or largely diluting a concentrated solution 
of potassium manganate, a brown precipitate of hydrated 
manganese dioxide is thrown down and a purple solution of 
potassium permanganate is left 

3K,MnQ 4 + 3H 2 = 2KMn0 4 + 4KOH + MnO,.H f O. 


A better yield of permanganate is obtained by blowing a 
current of carbon dioxide through the manganate solution, 
to combine with the alkali liberated. Concentration of the 
filtered solution yields dark purple crystals of the perman- 
ganate, isomorphous with those of potassium perchlorate, 
KC10 4 . Potassium permanganate is a powerful oxidizing 
agent. On heating it splits up into potassium manganate, 
manganese dioxide, and oxygen 

2KMn0 4 = K 2 Mn0 4 + MnO 2 + O 2 . 

It is usually employed for oxidizing purposes dissolved in 
water. The reaction follows a different course in alkaline 
solution from that in acid solution 

A. Alkaline Solutions. On boiling with an alkali solution, 
in presence of a substance that can be oxidized, the per- 
manganate is first reduced to manganate 

4KMn0 4 + 4KOH = 4K 2 Mn0 4 + 2H 2 + O t , 

and the manganate is then further reduced to manganese 
dioxide which is precipitated 

2K 2 Mn0 4 + 2H 2 = 2Mn0 2 + 4KOH + 0,. 

It follows that 2 molecules of potassium permanganate in 
alkaline solution yield 3 atoms of oxygen available for oxidation. 
Alkaline permanganate is used for oxidation purposes 
chiefly in organic chemistry ; the acid solution is generally 
employed in inorganic work. 

B. Acid Solutions. If we write the formula for two mole- 
cules of potassium permanganate, 2KMn0 4 , in the form 
K 2 O.2Mn0.50, we shall see that 2 molecules of the salt in 
acid solution are capable of yielding 5 atoms of oxygen available 
for oxidation. Sulphuric acid is used, since hydrochloric 
acid would be oxidized 

2KMn0 4 + 16HC1 = 2KC1 + 2MnCl a + 8H 2 + 5C1 2 . 

In the presence of sulphuric acid, then, potassium per- 
manganate is capable of acting in the following way, upon 
an oxidizable substance 

2KMa0 4 + 3H 2 S0 4 = K 2 S0 4 4- 2MnS0 4 + 3H 2 O + 5O. 


In acid solution, therefore, 316 grams of potassium per- 
manganate yield 80 grams of available oxygen. The equiva- 
lent of the salt in grams is thus 31-6 gms. ; ^ hence N/10 
potassium permanganate contains 3-16 gms. per litre. 

Standard potassium permanganate in the presence of 
sulphuric acid is one of the most valuable volumetric solu- 
tions. It is used for the estimation of hydrogen peroxide, 
ferrous iron, oxalic acid (on heating to 70), nitrites, and 
other oxidizable chemicals. The reactions which occur in 
the above examples are as follows 

(i) 50 + 5H 2 2 = 5H 2 + 50 2 . 

(ii) 50 + 10FeS0 4 + 5H 2 S0 4 = 5Fe 2 (S0 4 ), + 5H 2 O. 

(iii) 50 + 5H 2 C 2 4 = 10C0 2 + 5H 2 0. 

(iv) 50 -f 5HNO t - 5HN0 8 . 

The 50 in these equations is formed according to the equa- 
tion above. 

Sodium permanganate is the active agent in Candy's Fluid. 
The disinfecting and deodorizing properties of permanganates 
are due to the readiness with which these salts give up their 


1. What do you think of the position of manganese in the Periodic 
System ? Make a considered statement as to the general success or 
failure of the Periodic System in its grouping of the elements. 

2. Describe the preparation and properties of potassium per- 
manganate. For vi hat purposes is this substance employed in the 
laboratory ? 


THE " TEANSITION ELEMENTS." Iron, Cobalt, Nickel ; 

Ruthenium, Rhodium, Palladium ; Osmium, Iridium, 

The transition elements fall into three groups, indicated 
by semi-colons in the above list. The members of each 
group are very similar to one another, and have nearly the 
same atomic weights : thus, in round numbers 

(i) Fe = 56 (ii) Ru = 102 (iii) Os = 191 

Co = 59 Rh = 103 Ir = 193 

Ni = 58 Pd = 107 Pt - 195 

Groups (ii) and (iii) together are called the platinum metals. 
They are precious metals, extremely resistant to the action 
of acids and alkalis ; osmium, for example, will not dissolve 
even in aqua regia. The use of platinum in the laboratory 
is well known. 

The elements of group (i), iron, cobalt and nickel, are all 



Group in Periodic System : VIII ; Symbol : Fe ; Valency : 
2 and 3; Atomic Weight : 55-84; Melting Point : 1,500; 
Specific Gravity : 7-84. 

History. In spite of the fact that it is difficult to extract 
iron from its ores, this metal has been known for some 5,000 
years. In very early days it was more or less of a curiosity, 



but later became widely used for making swords and other 
weapons of war, and was supposed by the alchemists to be 
astrologically connected with the planet Mars, <? . About 
100,000 talents (2,500 tons) of iron were used in the construc- 
tion of SOLOMON'S temple (circa 1000 B.C.), and NEBUCHAD- 
REZZAR is said to have carried off into captivity in 604 B.C. 
a thousand blacksmiths from Damascus. Iron was worked 
in India from about 1000 B.C. and Indian steel was taken 
by sea to Oman and the Yemen where it was used by the 
Arabs for making fine blades. 

In the Old Testament iron is called barzil (from the ancient 
Assyrian barza), and the name " Berzelaios " in JOSEPHUS 
apparently means " the man of iron " an appropriate name 
for the chemist who first accurately determined the atomic 
weight of iron, BARON BERZELIUS ! Iron is mentioned also 
in the Qur'dn, Chapter 57, v. 25, " ALLAH sent down iron, 
in which is both keen violence and advantages to men," whence 
the names " The Violent " and " The Useful " often given to 
the metal by the chemists of Islam, e.g., ABU'L-QASIM AL- 
1 IRAQI (Book of the Mighty Secret of the Most Noble Stone], 
" Iron is hot and dry, related essentially to Marrikh (Mars), 
ruler of the Fifth Heaven. Its symbol is $ , and alternative 
names for it are ike. Violent, Death, the Useful, the Black 
Slave, and Persian earth." 

" GEBER " (thirteenth century) says, " Mars or Iron is a 
Metallick Body, very livid, a little red, partaking of Whitenefs, 
not pure, Jujtaining Ignition, fujible with no right fujion, 
under the Hammer exten/ible, and founding much. . . . 
Mars, among all the Bodies, is of leajt Perfection in Tranf- 
mutation, to be handled mojt difficult, and of exceeding long 

Iron is the Scandinavian and Anglo-Saxon name for the 

Occurrence. Iron forms about 4 per cent, of the earth's 
crust. It is found in small quantities free in nature, especially 
in meteorites, where it is usually associated with nickel. It 
occurs much more abundantly in the form of its oxides, 
sulphide, and carbonate 

IRON 545 

Oxides. Hcematite, Fe 2 8 ; magnetite, Fe 3 O 4 ; limonite, 
2Fe 2 O 3 .3H 2 O. 

Sulphide. Iran pyrites and marcasite, FeS 2 . 

Carbonate. Spathic iron ore or siderite, FeC0 8 . 

Iron is a constituent of the red colouring-matter of the 
blood (hcemoglobin) ; it is also essential to the life of green 
plants, although it is not present in chlorophyll. 

Extraction. The chief ores of iron are haematite (Lanca- 
shire, Belgium, North America), magnetite (Scandinavia, 
Germany, North America), limonite (South Wales), and 
spathic iron ore (mixed with clay as clay ironstone, in many 
parts of Great Britain). Iron pyrites is used chiefly as a 
source of sulphur. 

Iron is prepared from its ores by reducing them with carbon 
monoxide in a blast furnace 

Fe 2 O 8 + SCO ^=z 2Fe + 3CO 2 . 

The ore is first roasted to drive off any water present, to 
convert carbonate into oxide, to oxidize any sulphide to 
oxide, and to convert ferrous oxide (FeO) to ferric oxide 
(Fe 2 O 3 ), according to the ore used. 

The roasted ore is mixed with coke and limestone and the 
mixture introduced into the blast furnace through a cup 
and cone hopper at the top (see Fig. 118). The blast furnace 
is an iron tower lined with fire-resisting material, and is 
usually about 70-100 feet in height ; it is maintained at a 
temperature of from 700 at the bottom to over 1,500 at 
the top, chiefly by the heat evolved in the reactions. Hot 
air (700) is blown in at the bottom of the furnace through 
water- jacketed pipes called tuyeres (" twyers ") ; this oxidizes 
the adjacent coke into carbon monoxide which then reacts 
with the ferric oxide according to the equation given above. 
The reduced iron melts and sinks to the bottom of the furnace, 
accompanied by a molten slag of calcium aluminium silicate, 
formed from the clay and silica of the ores and the lime 
produced by decomposition of the added limestone. In this 
way, all the ore is converted into substances that can be 
removed from the furnace (namely, iron, slag, and carbon 




dioxide and monoxide), and the process is therefore con- 
tinuous. Fresh material is fed in from time to time and 
the furnace kept in operation until it is worn out,, An average 
furnace can produce about 500-600 tons of pig-iron a day. 

Reduction of 
Oxide to metal 


Carbon Monoxide 

Hot blast 

Waste gases used 
for heating blast 
and driving gas 

FIG. 118. Blast Furnace. 

The gases that escape from the furnace contain about 
GO per cent, of nitrogen, 12 per cent, carbon dioxide, and 24 
per cent, carbon monoxide. They are passed through flues 
into a gas -cleaner to remove the dust and are then burnt in 

IRON 547 

a brickwork " Cowper stove." When this is sufficiently 
heated the gases are turned into a, second Cowper stove while 
air is passed through the first. The hot air is then delivered 
through the tuyeres into the blast furnace. In this way 
one of the Cowper stoves is being heated while the other one 
is giving up its heat to the air-blast. This use of the " waste " 
gases effects a considerable saving of fuel. 

The dust which collects in the flues contains potassium 
salts, often in sufficient quantity to make their extraction 
commercially profitable. 

The molten iron that flows to the well of the furnace is 
run into moulds where it solidifies in blocks called pig iron 
or cast iron. It is very impure, containing as much as 4-5 
per cent, of carbon and 1 to 2 per cent, of silicon, as well as 
smaller quantities of phosphorus, sulphur and other elements. 
Cast iron is brittle and cannot be welded ; its use is therefore 

Wrought iron is made from pig iron by burning away the 
carbon, phosphorus, sulphur and silicon. The pig iron is 
melted in a reverberatory furnace lined with haematite 
(Fe 2 3 ), when the impurities are oxidized partly to gas 
(CO) and partly to a slag (Si0 2 , P 2 5 , etc.). The iron then 
begins to go " pasty," since the pure metal has a higher 
melting-point than that of the crude pig iron. The pasty 
iron is stirred and finally removed in large lumps, which are 
hammered with a steam hammer to remove the slag and 
then rolled into sheets. Further hammering and rolling 
produce the characteristic fibrous structure of wrought iron. 

Wrought iron is the purest form of commercial iron, con- 
taining about 99-99-5 per cent. Fe. It is very tough, has a 
high tensile strength, and can be welded (i.e. two pieces of it, 
when red-hot, can be joined together by hammering). It is used 
chiefly by blacksmiths and for the cores of electro -magnets. 

Steel is a mixture of iron and iron carbide or cementite, Fe 3 C. J 
The carbon content is from 0-2 to 1-5 per cent. Steel is now 
generally prepared by the SIEMENS -MARTIN open-hearth 

1 Hard steels are generally regarded as solid solutions of carbon in 



process, but the cementation and BESSEMER processes are still 
occasionally employed and will therefore be described here, 
(i) Cementation Process. Bars of wrought iron are packed 
with charcoal and heated to about 1,000 for 11-14 days. 
Absorption of charcoal by the iron gradually takes place and 
the product is called " blister-steel." In case-hardening the 
iron is removed when only the surface layers of the metal 

Basic Lining 

Supporting Ring 



FIG. 119. Bessemer Converter. 

have been converted into steel. Crucible steel (used for tools) 
is made by melting cementation steel in graphite crucibles, 
(ii) The BESSEMER Process. Cast iron is melted in a 
Bessemer converter (Fig. 119), and air is then blown through 
the molten metal to oxidize the carbon, phosphorus and other 
impurities. The converter is lined with a mixture of lime and 
magnesia (made by heating dolomite, CaC0 8 .MgC0 8 ), which 
absorbs the phosphorus pentoxide, forming Thomas slag or 
basic slag (crude calcium phosphate). The pure molten iron 
left in the converter is then mixed with a suitable weight of 



spiegeleisen (an alloy of iron, manganese and carbon) and 
thus converted into steel. 

The removal of phosphorus is extremely important, since 
the presence of phosphorus in steel renders the metal very 
brittle. Use of a basic lining to the converter was suggested 
by THOMAS and GILCHRIST (1879) and was very successful, 
since it not only efficiently removes the phosphorus but also 
yields a by-product (the " basic slag ") which is an excellent 

i^ Converter in Position 
* for Pouring 

Ladle in Position 
for Filling Ingots. 

FIG. 120. General Arrangement of Bessemer Plant. 

fertilizer for grass-lands. The general arrangement of the 
plant is shown in Fig. 120. 

(iii) Siemens -Martin Process. Most of the steel manu- 
factured at the present day is made by the Siemens-Martin 
process. A mixture of cast iron, scrap iron and haematite is 
placed on an open hearth and heated by producer-gas (p. 301) 
or by blast-furnace gases. The hearth is lined with calcined 
dolomite to remove phosphorus, as in the Bessemer process. 
By suitably adjusting the proportions of the cast iron and 


haematite, the carbon content of the product may be regulated 
as desired ; the molten steel is run off when ready by tilting 
the hearth. 

The properties of steel depend upon (a) its carbon content s , 
(b) its temper. 

(a) Carbon content. This varies from 0-2 to 1*5 per cent. 
The hardness of steel increases as the percentage of carbon in 
it rises ; steel which contains but little carbon is called " mild 

(b) Temper. If steel is heated to a high temperature and 
then suddenly cooled by being plunged into water it becomes 
extremely hard and brittle and is known as quenched steel. 
When quenched steel is carefully heated it becomes less hard 
but tougher, the extent of the change depending on the length 
of the heating and on the temperature. A rough indication 
of the temperature of the heated steel is given by the colour 
of the film of oxide formed on the surface, which varies from 
yellow through brown to blue as the temperature rises. Steel 
treated in this way is said to be tempered. The temperature 
at which the film of oxide is pale yellow is 200-230 ; the blue 
colour is produced at 290-300. Steel for cutting blades and 
tools is usually tempered yellow, while springs and chisels are 
generally tempered blue. 

Various forms of steel, suitable for a large number of different 
purposes, are made by the admixture of small quantities of 
other elements, e.g. nickel, chromium, tungsten, vanadium, 
manganese, molybdenum, cobalt, titanium and silicon. 
Chromium produces a harder steel, nickel a more elastic, 
molybdenum a tougher, and so on. Details of these special 
steels may be found in the handbooks of metallurgy. Nickel 
steels, containing 1-5 per cent, of nickel and 0-25-0-45 per cent, 
of carbon, and nickel-chromium steels (1-4-5 per cent, nickel, 
1-2 per cent, chromium) are used for propeller shafts, piston 
rods, crankshafts, steering gears, axles, brake rods, etc. Stain- 
less steel contains about 12-15 per cent, chromium. Silicon 
fteel, or silico-manganese steels (1-2 per cent, silicon, 0-4-1 
per cent, manganese, 0-40-0-65 per cent carbon) are used 
especially for making springs. Manganese steels containing 

IRON 551 

12-14 per cent, manganese are used for tramways points and 
crossing-rails, as they are very hard and tough. 

Properties. Iron exists in four allotropic modifications, 
very similar to one another in general chemical properties. 
The form stable at ordinary temperatures is called a-iron or 
" ferrite " ; at 766 this changes into /8-iron, which is stable 
from 766 to 895 ; at 895 /ff-iron changes into y-iron and 
this is converted into <5-iron at 1,400. Quenched steel is a 
solution of cementite (Fe 3 C) in y-iron ; the various tempered 
steels consist of heterogeneous mixtures of a-iron, or " ferrite," 
with cementite. 

Pure iron is a rather soft, lustrous, white, magnetic metal, 
with specific gravity 7-84, melting-point 1,500, and boiling- 
point 2,450. It dissolves in dilute sulphuric and hydro- 
chloric acids with evolution of hydrogen ; concentrated 
sulphuric acid dissolves it on heating, with formation of 
sulphur dioxide. It is dissolved by cold dilute nitric acid, 
forming ferrous nitrate and ammonium nitrate ; concentrated 
nitric acid, however, renders it passive. In the " passive " 
state iron will not dissolve even in dilute acids ; the phenome- 
non is supposed to be due to the formation of a coherent 
protective film of oxide on the surface of the metal, and i& 
shown by certain other metals, e.g. chromium. Passivity i& 
removed by heating the metal in hydrogen or by rubbing with 
sand-paper. Further information may be obtained from a 
paper by HEATHCOTE (Jour. Soc. Chem. Ind., 26, 899, 1907). 

In moist atmospheric air ordinary iron rapidly rusts to a 
brown powder which has the composition Fe 2 3 .H 2 or 
2Fe 2 O 3 .3H 2 0. It is said that absolutely pure iron will not 
rust. The mechanism of rust -formation is a problem that 
has attracted the attention of a great number of chemists and 
cannot even yet be regarded as satisfactorily solved. It is 
probable that when ordinary iron rusts under normal atmo- 
spheric conditions ferrous bicarbonate, Fe(HC0 3 ) a , is first 
formed ; this dissolves in the water present and is oxidized 
by the dissolved oxygen in the water to a hydrated ferric 
oxide, with liberation of carbon dioxide. Other chemists 
regard the rusting of iron as essentially an electrochemical 


phenomenon. On this theory, different bits of the piece of 
iron, varying in physical or chemical condition (e.g. local 
compression or strain, presence of small quantities of im- 
purity) may act as the positive and negative plates of a battery 
when covered by air or impure water. Corrosion would 
then take place at those parts acting as negative poles, i.e. 
where the concentration of oxygen is least, as in cracks and 
crevices in the surface. It has been shown that the rate of 
corrosion is affected by many different factors, and the problem 
is an extremely complicated one. The student is recommended 
to read the following papers 

DUNSTAN, JOWBTT, AND GouLDiNG, Trans. Chem. Soc., 87, 
1,548 (1905) ; MOODY, ibid., 89, 720 (1906) ; WHITNEY, Jour. 
Amer. Chem. Soc., 25, 394 (1903) ; TILDEN, Trans. Chem. 
/Soc., 93, 1,356 (1908) ; LAMBERT AND THOMSON, ibid., 97, 
2,426 (1910) ; DUNSTAN AND HILL, ibid., 99, 1,835 (1910) ; 
LAMBERT, ibid., 101, 2,056 (1912) ; U. K. Evans, The Corrosion 
of Metals (London, Edward Arnold & Co., 1926). 

Iron will combine directly with many non-metallic elements 
if heated with them ; thus with sulphur it yields ferrous 
sulphide, in oxygen it burns brilliantly to magnetic oxide of 
iron, and with chlorine it gives ferric chloride. It does not 
form a hydride, but red-hot iron will allow hydrogen to pass 
freely through it. 

COMPOUNDS OF IRON. Iron forms three oxides and two 
series of salts 

FeO, ferrous oxide, to which correspond ferrous salts. 
Fe 3 4 , ferroso-ferric oxide or magnetic oxide of iron. No 

corresponding salts. 
Fe 2 3 , ferric oxide, to which correspond the ferric salts. 

Solutions of ferrous salts contain the bivalent ionF" ; those 
of ferric salts contain the tervalent ion Fe'". Ferrous salts 
are reducing agents since they are readily oxidized to the 
ferric state- 

Ferrous oxide, FeO. 

Ferrous oxide is obtained as a black powder by gently 
heating finely divided iron with nitrous oxide or by carefully 
reducing ferric oxide with hydrogen at 300. Mixed with iron 

IRON 553 

it can be made by igniting ferrous oxalate in absence of air, 
FeC 2 O 4 - FeO + CO + C0 2 . 

It takes fire spontaneously in the air, burning with incan- 
descence to ferric oxide, Fe 2 O 3 . It dissolves in acids, forming 
ferrous salts. 

Ferrous hydroxide, Fe(OH) 2 , is precipitated as a white 
gelatinous solid by addition of caustic soda solution to a solu- 
tion of a ferrous salt in absence of air. It rapidly oxidizes in 
the air, forming first the green compound Fe(OH) 2 .2Fe(OH) 8 
and finally brown ferric hydroxide Fe(OH) 3 . 

Ferrous chloride, FeCl 2 , is made in the anhydrous state 
by heating iron wire in a current of hydrochloric acid gas 
Fe + 2HC1 = FeCl 2 + H 2 . 

It is a white crystalline solid and can be vapourized 
unchanged. Vapour density determinations show that 
the molecules are at first Fe 2 Cl 4 ; on further heating these 
dissociate into FeCl 2 

Fe 2 Cl 4 ;=z 2FeCl,. 
The dissociation is complete at 1,500. 

The tetrahydrate, FeCl 2 .4H 2 O, is most conveniently made 
by dissolving iron in dilute hydrochloric acid and evaporating 
the solution. It forms pale-green crystals. 

Ferrous sulphate, or green vitriol, FeSO 4 .7H 2 O, may be 
made by dissolving iron in dilute sulphuric acid and evapor- 
ating the solution to crystallization. Commercially it is made 
by exposing marcasite, FeS 2 , to moist air 

2FeS 2 + 2H 2 O + 70 3 = 2FeS0 4 + 2H 2 SO 4 . 

The anhydrous salt may be obtained as a white powder by 
cautiously heating the hydrate in absence of air. 

Other hydrates are known, e.g. FeSO 4 .5H 2 O and 
FeSO 4 .H 2 O. Ferrous sulphate crystals are slowly oxidized 
in the air and therefore in volumetric analysis it is more con- 
venient to make a standard solution of ferrous iron from the 
double salt ferrous ammonium sulphate, 

FeS0 4 .(NH 4 ; 2 S0 4 .6H a O, 

which is much more stable. 



When strongly heated ferrous sulphate decomposes, yield- 
ing oxides of sulphur, steam, and ferric oxide. If the gases are 
cooled in a receiver fuming sulphuric acid (pp. 468, 474) is 

The brown compound formed when nitric oxide dissolves in 
ferrous sulphate solution has been described previously (p. 363). 

Ferrous sulphate is largely used in making ink, which 
consists of a mixture of a dark-blue dye with tannin and 
ferrous sulphate solution. When fresh, writing in such an 
ink is blue, but goes black gradually on exposure to air owing 
to the formation of an oxidized iron-tannin compound. 

Ferrous sulphide, FeS, is made by heating iron and 
sulphur together or by adding sodium sulphide to ferrous 
sulphate solution 

FeS0 4 + Na 2 S - FeS + + Na 2 S0 4 . 

It is a black solid that melts at about 1,200 and can easily 
be cast in sticks. It is used for the preparation of sulphuretted 
hydrogen. The commercial substance always contains un- 
combined iron, to which it probably owes its black colour ; it 
is said that the pure substance forms brownish-yellow crystals. 

Ferrous carbonate, FeC0 3 , occurs naturally as spathic 
iron ore and, mixed with clay, in clay ironstone. It comes 
down as a white precipitate on addition of sodium carbonate 
to ferrous sulphate solution ; on exposure to air it oxidizes, 
becoming first green and then brown. It will dissolve if 
suspended in water through which a current of carbon dioxide 
is blown, owing to the formation of soluble ferrous bicarbonate 
FeC0 3 + H 2 + C0 2 = Fe(HC0 3 ) 2 . 

Like the normal carbonate, the bicarbonate is quickly 
oxidized in air. 
Ferroso -ferric oxide, Fe 3 O 4 . 

Ferroso -ferric oxide occurs naturally as magnetite or load- 
stone. It is formed when iron burns in air or oxygen, and 
when iron is heated in a current of steam 
3Fe + 4H 2 = Fe 8 4 + 4H 2 . 

Chemically it behaves as a loose " molecular compound " 
of ferrous oxide and ferric oxide. The fused oxide is not 


attacked by acids, oxygen, or chlorine and is therefore some- 
times used for making anodes for electrolytic processes. 
Ferric oxide, Fe 2 O a . 

Ferric oxide occurs in the anhydrous state as the important 
mineral hcematite, and in the hydrated state as liinonite, 
2Fe 2 3 . 3H 2 0. It can be made by igniting ferric hydroxide 
or ferrous sulphate ; prepared from the latter it is known as 
jeweller's rouge or colcolhar and is used as a polishing powder 
and as a red pigment. 

Ferric hydroxide, Fe(OH) 3 , is formed as a reddish-brown 
precipitate by adding caustic alkali solution or ammonia to 
a solution of a ferric salt. It is a very weak base, and is 
soluble in acids, forming ferric salts. It is so weak a base 
that it will not form a carbonate ; hence it is precipitated 
when a solution of sodium carbonate is added to a solution 
of a ferric salt. 

Ferric chloride, FeCl 3 , is formed as a black scaly crystal- 
line solid by heating iron in a stream of chlorine. It can be 
vapourized unchanged, and vapour density determinations 
show that at 445 the molecules are Fe 2 Clj. On further 
heating dissociation takes place, first into 2FeCl 3 and finally 
into ferrous chloride and chlorine. 

Fe 2 Cl 6 ^= 2FeCl 3 ^= 2FeCl 2 + C! 2 . 

Ferric chloride forms many hydrates, all of which are 
yellow in colour 

FeCl 3 .6H 2 O (ordinary "ferric chloride crystals"). 
Fed 3 .3|H 2 O. FeCl 3 .2iH 2 0. FeCl 3 .2H 2 O. 

Aqueous solutions of ferric chloride are strongly acid, 
owing to hydrolysis. Ferric hydroxide is a weak base and 
is therefore formed from the hydroxyl ions of the water and 
the ferric ions of the ferric chloride. Each hydroxyl ion 
removed in this way leaves a free hydrogen ion in solution and 
hence the solution becomes acid 

(i) H 2 O ^n IT + OH'. 
(ii) FeCl 3 ^r Fe* " + 3d'. 
(iii) Fe- + 30H' ^r Fe(OH) . 



FeCl, + 3H 2 = Fe(OH), + 3HC1. 


If a solution of ferric chloride is dialysed the hydrochloric 
acid passes through the membrane and a colloidal solution of 
ferric hydroxide is left in the dialyser. This solution is used 
in medicine under the name of " dialysed iron." 

Ferric sulphate, Fe 2 (S0 4 ) 8 is usually met with in the form 
of its double salt with ammonium sulphate, ferric alum, 
(NH 4 ) a S0 4 .Fe 2 (SO 4 ) 3 .24H 2 0. This salt is isomorphous 
with ordinary potash alum and generally has a pale violet 
colour. Potassium ferric alum, K 2 S0 4 .Fe 2 (S0 4 ) 3 .24H a O, 
is very similar. 

Ferric sulphate is formed by heating ferrous sulphate with 
a mixture of sulphuric and nitric acids 

6FeS0 4 + 3H 2 S0 4 + 2HN0 3 = 3Fe 2 (S0 4 ) 3 + 4H 2 + 2NO. 

The anhydrous salt is a white or yellowish powder. 

Ferric sulphide, Fe 2 S a , is a black substance that may be 
made by passing sulphuretted hydrogen through a suspension 
of ferric hydroxide in water 

2Fe(OH) 3 + 3H 2 S = Fe 2 S 3 + 3H 2 0, 

or by adding a solution of a ferric salt to an excess of 
ammonium sulphide solution. In presence of excess of sul- 
phuretted hydrogen it is converted into iron disulphide, 
FeS a . 

Iron disulphide, FeS 2 , exists naturally in two different 
crystalline forms, iron pyrites and marcasite. Marcasite is 
much more readily oxidized by moist air than is pyrites. 

Iron carbonyls. When iron is heated in carbon monoxide 
it forms a compound, Fe(CO) 6 , known as iron pentacarbonyl. 
This is a yellow liquid boiling at 103 and freezing at 21. 
In the presence of light it gradually decomposes into another 
carbonyl of iron, Fe 2 (CO) 9 

2Fe(CO) 5 = Fe 2 (CO) 9 + CO. 
Iron tetracarbonyl, Fe(CO) 4 , is also known. 

IRON 557 

All the carbonyls of iron are decomposed at a fairly high 
temperature into iron and carbon monoxide. 

Organic Compounds of Iron. Iron forms many impor- 
tant organic compounds, the chief of which are potassium 
ferrocyanide, K 4 Fe(CN) 6 , potassium ferricyanide, K 3 Fe(CN) 6 , 
and sodium nitroprusside, Na 2 Fe(NO)(CN) 5 . For the pre- 
paration and properties of these compounds, see p. 215, and 
text-books of organic chemistry. Consult also FENTON'S 
Notes on Qualitative Analysis for the use of ferrous salts as 
catalysts in certain organic preparations, and for the 
detection of iron. 

Estimation of Iron.- Ferrous iron is usually estimated 
by addition of sulphuric acid and titration with standard 
potassium permanganate or potassium dichromate. Ferric 
iron may be estimated by reduction to the ferrous state (e.g. 
with zinc and acid), followed by titration with permanganate, 
or preferably by direct titration with standard titanous 
chloride, using ammonium thiocyanate as indicator. See 
BERRY, Volumetric Analysis, and KNECHT and HIBBERD, 
New Reduction Methods in Volumetric Analysis. 

Titration of ferrous sulphate with (i) Permanganate 

10FeS0 6 + 8H 2 S0 4 + 2KMn0 4 

= K 2 S0 4 + 2MnS0 4 + 5Fe. 2 (S0 4 ) 3 + 8H 2 0. 

(i.e. 316 grams KMn0 4 = 560 grams Fe). Indicator : None. 
End-point : Permanent pink colour, 
(ii) Dichromate 

6FeSO 4 + 7H 2 S0 4 + K 2 Cr 2 7 

= K 2 S0 4 + Cr 2 (S0 4 ) 3 + 3Fe 2 (S0 4 ) 3 + 7H 2 O. 
(i.e. 2934 grams K 2 Cr 2 7 == 336 grams Fe). Indicator : 
Potassium ferricyanide solution, externally, on white tile. 
End-point : Red-brown colouration. 

Titration of ferric chloride with titanous chloride 

FeCl 3 + TiCl 3 = FeCl a + TiCl 4 . 

The titanous chloride solution is standardized against a 
standard ferric solution. Indicator : Ammonium (or potas 
sium) thiocyanate. End-point : Blood-red colour vanishes. 



Group in Periodic System : VIII ; Symbol : Co ; Valency : 
2 and 3 ; Atomic Weight : 58-94 ; Melting -Point : 1450 ; 
Specific Gravity : 8-8. 

Occurrence. It is said that the word cobalt is etymolo- 
gically connected with kobold, the name given to certain 
gnomes who lived in copper mines in the Hartz Mountains and 
took a spiteful glee in deceiving men by exposing ores which 
looked like copper ores but from which no copper could be 
extracted. These ores are now known by the name of 
smaltine (CoAs 2 ) and cobaltine (CoAs 2 . CoS 2 ), and are obtained 
at the present day chiefly from North America (Ontario, 
etc.). The word kibaltu occurs among the names of other 
minerals on an ancient Assyrian tablet dealing with glass- 
making, but it is uncertain whether the word refers to a cobalt 
ore or not. 

Extraction. The cobalt ores are first roasted in a current 
of air, which removes the arsenic and leaves the cobalt in the 
form of oxide. This is then reduced with charcoal or hydro- 
gen. As the cobalt ores are seldom pure, special methods of 
separation are generally necessary. 

Properties. Cobalt is a silvery metal with a reddish 
lustre. Its specific gravity is 8-8, and it melts at about 1450. 
It is unchanged on exposure to air at ordinary temperatures, 
but oxidizes gradually when heated. It is readily soluble in 
nitric acid but dissolves only very slowly in hydrochloric or 
sulphuric acids. It is magnetic but much less so than iron, 
although it retains its magnetic properties up to a temperature 
of over 1,000 C. Finely divided cobalt has the power of 
absorbing many times its own volume of hydrogen. 

COMPOUNDS OF COBALT. Cobalt forms two main oxides 
and two corresponding series of salts. 

Cobaltous oxide, CoO, is a greenish-brown powder ob- 
tained by gently heating cobaltous hydroxide or carbonate 
out of contact with air. Cobaltous salts are pink in solution 
or when they contain water of crystallization, but the anhy- 
drous salts are blue. 


Cobaltous chloride, CoCl 2 .6H 2 O, is obtained by dissolving 
cobaltous oxide or carbonate, or the metal itself, in dilute 
hydrochloric acid. On evaporating the solution, red crystals 
of cobaltous chloride hexahydrate separate out. Other 
hydrates are known. The anhydrous salt is blue, and this 
fact has been made use of in the preparation of " invisible " 
or " sympathetic " ink. When a dilute solution of cobaltous 
chloride is used for writing, the characters are invisible, but 
on heating the paper over a lamp or before the fire the hy~ 
drated salt loses its water of crystallization and the blue 
anhydrous salt is formed, the characters becoming apparent 
at once. If the paper be now breathed on, or exposed to 
damp air, the writing fades again. Paper soaked in cobaltous 
chloride solution and then dried is sometimes used as a rough 
weather-glass, as when the air is very hot and dry the paper 
goes blue, becoming pink again on the approach of rain ; the 
paper may also be used for showing the transpiration from the 
under surface of dorsi ventral leaves. 

An interesting observation is that a solution of cobaltous 
chloride goes blue (a) on heating, if sufficiently concentrated ; 
(6) on addition of strong hydrochloric acid ; (c) on addition 
of strong sulphuric acid ; and (d) on addition of alcohol. 
These facts may be explained by assuming a diminution in 
the ionization of the salt brought about under the above con- 
ditions ; it is considered that the cobaltous ion is pink, while 
the un-ionized molecules CoCl 2 are blue. It is possible, 
however, that a complex anion of a deep blue colour is formed ; 
certainly cobalt forms many complex compounds, the colours 
of which are very varied. 

Cobaltous sulphide is the black precipitate obtained in 
Group IV of the Analysis Tables. 

Cobaltous carbonate is a pink substance precipitated 
when sodium carbonate solution is added to a solution of 
cobaltous chloride. Solutions of cobaltous salts yield a 
precipitate of red cobaltous hydroxide when treated with 
caustic potash solution. The precipitate is soluble ID 

Cobalt silicate is deep blue in colour and is used in the 


preparation of blue glass, enamel and porcelain. The paint 
" cobalt blue " is the powdered silicate. 

Cobaltic oxide, Co 2 8 , is obtained as a black powder by 
carefully heating cobaltous nitrate ; further heating results 
in the formation of cobalto-cobaltic oxide, Co 3 4 . Co 2 8 may 
also be formed by adding bleaching-powder solution to a 
solution of cobaltous chloride. It dissolves in acids, but the 
cobaltic salts so formed are unstable and readily decompose, 
forming cobaltous salts with liberation of oxygen, etc. 

Complex Compounds of Cobalt. Cobalt forms a great 
number of complex compounds, which have recently become 
of importance in connection with the work of WERNER and 
others on valency. The compounds formed by cobalt salts 
and ammonia are called the cobaltammines ; several series of 
these are known, and they are classified, according to their 
colour, as roseocobalt ammines, purpureo-, xantho-, croceo-, and 
flavo- cobalt ammines, etc. Preparation of some of these salts 
is an interesting occupation for spare moments in the labora- 

Potassium cobalticyanide, K 3 Co(CN) 6 , is a pale yellow crystal- 
line solid, resembling potassium ferricyanide, K 3 Fe(CN) 6 , 
with which indeed it is isomorphous. Potassium cobalto- 
cyanide, K 4 Co(CN) 6 , is also known. 

Potassium cobaUinitrite, K 3 Co(N0 2 ) 6 .l7fH 2 0, is obtained as 
a yellow precipitate by adding potassium nitrite to a solution 
of a cobaltous salt containing a little acetic acid. It is 
slightly soluble in water. Nickel does not form an analogous 
compound, so the two metals may be separated in this way. 


Group in Periodic System : VIII ; Symbol : Ni ; Valency : 
2 and 3; Atomic Weight: 58*68; Melting Point : 1452; 
Specific Gravity : 8- 9. 

History. Certain minerals that resembled copper ore 
but from which no copper was obtainable were called by the 
mediaeval German miners kupfernickel or " false copper." 
From these minerals KRONSTEDT in 1751 obtained nickel. It 


is to BERGMANN, however (1774), that we owe the first accur- 
ate description of nickel and its compounds. 

Occurrence. Nickel occurs in small quantities native, in 
meteorites, but is chiefly found as kupfernickeL NiAs. smaltine 
(Ni,Fe,Co)As 2 , nickel glance, NiAsS, nickel blende, NiS, pent- 
landite (Cu,Fe,Ni)S, and garnierite, a silicate of magnesium 
and nickel of complex composition. The chief deposits are at 
Sudbury (Ontario) and New Caledonia. 

Extraction. Nickel is chiefly extracted by MOND'S car- 
bonyl process. The ores are first roasted, when the sulphur 
is converted into sulphur dioxide and the nickel and other 
metals left as oxides. These are then heated in a stream of 
water-gas (CO + H 2 ), which reduces the nickel oxide to 
nickel ; the temperature, however, is adjusted so that the 
iron oxide is not reduced. 

The product is then placed in a tower (the " volatilizer ") 
and heated to a temperature of 80-100 in a current of carbon 
monoxide. This combines with the nickel to form a volatile 
substance, nickel carbonyl, Ni(CO) 4 , which is passed into 
another tower at a higher temperature (180), where it is 
decomposed into nickel and carbon monoxide. The carbon 
monoxide is led off and used again. 

Ni + 4CO ^z Ni(CO) 4 . 

The nickel so produced is of a high degree of purity (about 
99'8 per cent.). It can be further purified, if necessary, by 

Properties. Nickel is a greyish metal that can be 
polished till it looks like silver. Its specific gravity is 8-9 and 
its melting-point 1,500. It is stable in dry air and tarnishes 
only slowly in moist air ; it is therefore used for nickel- 
plating other metals, on which it may be deposited by electro- 
lysis. In chromium-plating, it is used to deposit the chromium 
on an underlying coating of nickel. In general and magnetic 
properties nickel resembles cobalt. 

It is a constituent of many alloys, used for coinage, etc. 
Nickel-steel is steel containing varying amounts of nickel 
(usually about 10 per cent, by weight ) and is used for armour- 


plates on account of its extraordinary toughness, which 
renders it less likely to develop serious cracks when struck by 
shells. % 

Nickel is also widely employed as a catalyst in both pure 
and applied chemistry and in industry (see p. 185, etc.). 


Nickel forms at least two oxides, NiO and Ni 2 8 . The 
former is basic and gives rise to the nickel salts, in which 
nickel is bivalent. The nickel ion Ni" is green, as are most 
nickel compounds. 

Nickel carbonyl, the compound used in the extraction of 
the metal, is a colourless liquid boiling at 43 C. It has the 
formula Ni(CO) 4 , as shown by vapour density determinations. 
On strongly heating it explodes, forming nickel, carbon, and 
carbon monoxide. It burns in air with a smoky flame. 

Nickel may be detected in margarine, fats, etc., where it is 
liable to occur occasionally (see hydrogen, p. 185), by means 

CH 8 C = N . OH 

of dimethylglyoxime, \ , with which it gives 

CH 3 C-N.OH 

a scarlet precipitate. (See FENTON'S Notes on Qualitative 


1. What do you know of the history of iron ? 

2. Describe the metallurgy of iron. 

3. What is steel ? How is it made ? 

4. How would you prepare the anhydrous chlorides of iron ? What 
are their properties ? 

6. Write a general account of the compounds of cobalt. 

6. How is nickel extracted from its ores ? 

7. How far do iron, nickel and cobalt resemble one another t 


Acheson, 280 

Agricola, 403 

Al-Andalusi, 4 

Albertus Magnus, 391 

Albright, 375 

Aldus, 399 

Al-'Iraqi, 4, 20, 228, 391, 544 

Al-Jildaki, 4, 19 

Al-Razi, 4 

Arnmon, Jupiter, 338 

Ampere, 528 

Andrews, 78 

Arabs, 2-5, 210, 259, 262, 268, 273, 

314, 321, 390, 496, 537, 544 
Aristotle, 2, 315, 317 
Armstrong, 441 
Arrhenius, 7, 129 
Aston, 174 
Avogadro, 35, 177, eto. 

Bacon, Roger, 5 
Baker, 263, 415 
Balard, 249, 510, 611, 518, 526, 


Barkla, 166 
Baron, 269 

Becher, 6, 308, 410, 454 
Becquerel, 8, 158 
Bergius, 287 

Bergmann, 117, 403, 561 
Berkeley, Earl of, 103 
Berry, 250, 557 
Berthelot, 8, 239, 477 
Berthollet, 118, 496 
Berzolius, 34, 60, 93, 117, 250, 309, 

Bessemer, 548-9 

Biltz, 461 

Birkeland, and Eyde, 355 

Black, 240, 303 

Blagdon, 106, 107 

Boehringer, 351 

Bohn, 211 

Bohr, 177 

Boltzmann, 70 

Bouasingault, 450 

Boyle, 5, 25, 27, 43, 373, 376 

Bradley, 313 

Bragg, Sir W. H., 163 

Bragg, W. L., 163 

Brand, 372 

Bray, 512 

Bredig, 145, 222 

Brodie, 287, 422 

Brook, van den, 166, 16J> 

Brown, J. Campbell, 8 

Bunsen, 57, 75, 244 

Burt, and Edgar, 440 

Burton, Sir R. F., 438 

Bussy, 240 

Butterfield, 292 

Caesar, 314 
Cannizzaro, 35, 151 
do Carli, 271 
Carlile, 93 
Carrara, 431 
Carre, 343 
Casciorolus, 251 
Cassius, 146, 237, 312, 320 
Castner, 194 
Castoro, 320 

Cavendish, 19, 179, 187, 353, 410. 




Chance, 204, 457 
Chancel, 378 
Charles, 25, 43 
Charles the Hammer, 5 
Chatelier, Le, 85-7, 123, 340, 462, 


Clark, 446 
Claus, 457 
Clausius, 70 
Clement, 468 
Clerk-Maxwell, 70 
Condy, 542 
Congreve, 378 
Conn, 513 
Courtois, 528 
Cowper, 547 
Crafts, and Friedel, 279 
Crawford, 250 

Crookes, Sir W., 159, 257, 372 
Cullis, 322 
Curie, M. and Madame, 8, 158 

Dalton, 7, 9, 10, 14-17, 33, 35, 60, 

177, 211, 414 
Darling, Mr. Justice, 13 
Davy, Sir H., 194, 211, 240, 244, 

250, 252, 269, 273, 360, 497, 512, 

513, 614, 528 
Deacon, 142 
Democritus, 2 
Desormes, 468 
Deville, 274, 367, 483 
De Vries, 100 
Dewar, Sir J., 80, 183, 293, 420, 


Dioscorides, 2, 258, 398 
Dippel, 374 
Dirac, 177 
Divers, 470 
D6bereiner, 150 
Drechsel, 299 
Dujardin, 455 

Dulong, and Petit, 54, 57, 91, 155 
Dumas, 44, 47, 440, 449, 450 
Dunstan, 552 
Dutrochet, 100 

Edgar, and Burt, 440 
Empedocles, 116 
Erdmann, 317, 624 
Eyde, and Birkeland, 355 

Faraday, 21, 94-7, 128, 342, 603, 

506, 512 
Fehling, 222 

Fenton, 143, 390, 397, 557, 562 
Ferdinand II of Bourbon, 469 
Fielding, 269 
Findlay, 147 
Fischer, 100 
Fischer, E., 7 
Follie, La, 468 
Frasch, 456 
Freund, Miss, 9 
Friedel, and Crafts, 279 

Gahn, 251, 373, 538 

Gaillard, 471 

Garbett, 468 

Gay-Lussac, 34, 249,269, 308, 470, 

492, 497, 511, 513, 528, 534 
Geber, 4, 218, 222, 228, 234, 259, 

262, 264, 314, 322, 328, 338, 347, 

352, 391, 544 
Gengembre, 37P 
Gooff roy, 177 
Gilchrist, 649 
Gill, 455 

Gill and Payen, 455 
Gladstone, 121 
Glauber, 208, 352, 496 
Glover, 471 
Goldschmidt, 514 
Gore, 400 
Gossage, 196, 202 
Goulding, 552 
Graham, 74, 144 
Greeks, 2 
Grew, 240 

Guldberg, and Waage, 120 
Guye, 521 

Haber, 341 
Hall, 274 
Halske, 423 
Hargreaves-Bird, 197 
Hartley, 103 
Hautefeuille, 183 
Heathcote, 551 
Henckel, 251 
Heraclitus, 2 
Hermes, 260 
Heroult, 274 
Hess, 89 



Hibberd, 557 
Hill, 552 
Hindus, 1 
Hippocrates, 2 
Hofmann, 49, 349 
Holker, 468 
Holmes, 381 
Homberg, 251, 269 
Homer, 420 
Hope, 250 
Hundeshagen, 528 

Ibn Arfa 'Ras, 4 
Israelites, 237 

Jabir ibn Hayyan, 4, 177, 468 
Josephus, 544 
Joule, 79 
Jowett, 552 
Jupiter Ammon, 338 

Kammerer, 527 
Kekule, 7 
Kelvin, Lord, 79 
Kingzett, 515 
Kipp, 304, 462 
Kipping, 314 
Klaproth, 250, 481 
Knecht, 557 
Krafft, 373 
Kronstadt, 660 
Kunckel, 373 

La Follie, 468 

Lambert, 322, 552 

Lamy, 455 

Langmuir, 8, 177, 185 

Laue, 161 

Lavoisier, 7, 179, 181, 211, 244, 

269, 283, 303, 308, 332, 353, 373, 

412-14, 454, 464, 496 
Leblanc, 200, 479, 500 
Le Chatelier, 85-7, 123, 306, 340, 

462, 472 

Leclanche, 347, 640 
Lehmann, 483 
Lemery, 269, 376, 403, 405 
Le Sueur, 198 
Lewis, 177 
Libavius, 320, 496 
Liebig, 7, 518 
Lipowitz, 404 

Lippmann, 8 
Lockyer, Sir N., 188 
Lowry, 8 

McBain, 184 

Mackenzie, 283 

Marggraf, 273 

Marsh, 397 

Martin, 547, 549 

Mary the Copt, 142 

Mayow, 117 

Mazzetti, 271 

Mendeleeff, 151-7 

Mercury (Hermes), 260 

Meyer, E. von, 8 

Meyer, L., 155 

Meyer, Victor, 44, 48-50, 67, 345 

Michael, 513 

Michel, 455 

Millon, 264, 533 

Mills, 277, 347 

Mitscherlich, 58 

Moissan, 244, 284, 286, 483, 490, 

492, 493, 495, 638 
Moles, 521 
Molinari, 469 
Mond, 561 
Moody, 552 
Morley, 440 
Morse, 102 
Moseley, 166 
Moses, 237 
Muspratt, 201, 515 
Mutanabbi, 398 

Nebuchadrezzar, 644 
Nernst, 7, 242 
Nessler, 266, 350 
Newlands, 151 
Newton, 9, 117, 404 
Nicholson, 93 
Nollet, Abbe, 100 

Odling, 249 
Onnes, 78, 189 
Orleans, Duke of, 200 
Ostwald, 7, 133, 141 

Paracelsus, 4, 6, 251, 273, 397, 403 
Parker, 375 
Parkes, 229 
Parrish, 385 



Parrot, 100, 101 

Parsons, Sir C., 285 

Partington, 394 

Pasteur, 7 

Pattinson, 229 

Payen, and Gill, 455 

Peachey, 321 

Pebal, 514 

Peligot, 489 

Petit, and Dulong, 54, 57, 91, 155 

Pettenkofer, 150, 452 

Pfeffer, 101 

Pharaoh, 266 

Phillips, 473 

Planck, 91 

Plato, 2 

Plato, and Ruff, 244 

Pliny, 259, 314, 338, 637 

Pope, Sir W. J., 7, 321 

Pott 273 

Priestley, 6, 261, 264, 332, 338, 

360, 363, 408, 412-4, 464, 496 
Pring, 269 
Pringsheim, 100 
Ptolemy, King, 338 
Pythagoras, 2 

Ramsay, Sir W., 7, 187, 188 

Raoult, 7, 106, 109, 110 

Raschig, 470 

Rayleigh, Lord, 44, 187 

Readman, 375 

Regnault, 43, 44 

Retgers, 377 

Reynolds, 470 

Richards, 527 

Robinson, 375 

Roebuck, 468 

Rontgen, 158, 169 

Roozeboom, 183 

Rose, 404 

Ruff, and Plato, 244 

Rulandus, 403 

Rutherford, 332 

Rutherford, SirE., 8, 165, 168, 174 

Rydberg, 166 

Sabatier, 185, 292, 367 
Sargon, 239 
Sauria, 378 
Schaffner, 324 

Scheele, 251, 308, 332, 373, 394, 

408, 492, 496, 538 
Schonbem, 421 
Schoolmen, 32 - 
Schroder-Grillo, 475 
Schrodinger, 177 
Schrotter, 377 
Schweitzer, 225 
Scott, 520 

Senderens, 185, 292, 367 
Shelburne, Lord, 413 
Shenstone, 422, 499 
Shipton, Mother, 9 
Siemens, 422, 547, 649 
Smith, 515 
Soddy, 8, 168 
Solomon, 544 
Solvay, 204-5 
Soret, 425, 426 
Stahl, 6, 410, 454 
Stas, 461, 521 
Stewart, 172 
Stopes, 289 
Strohmeyer, 258 
Strutt, 337 
Sutton, 448 

Tachenius, 308 

Tafel, 351 

Taix-Aycard, 469 

Tammann, 109 

Taylor, 470 

Tennant, 283 

Thenard, 269, 308, 431, 492, 497 

Thorn, 205 

Thomas, 548, 549 

Thomson, 652 

Thomson, G., 303 

Thomson, Sir J. J., 8, 171 

Thomson, Sir W., 79 

Thorpe, 280 

Tilden, Sir W., 552 

Traube, 101 

Troost, 183 

Valentine, Basil, 5, 398, 403, 492 

van den Broek, 166, 169 

van der Waals, 77 

van Helmont, 303 

van Marum, 420 

van't Hoff, 7, 104, 127 



Vauquelin, 483 
Venus, 218 

Waage, and Guldberg, 120 
Wackenroder, 478 
Ward, 468 
Warth, 270 
Weintraub, 269, 270 
Weldon, 501, 540 
Werner, 486, 560 
Western, 263 
Whitney, 552 

Wilhelmy, 120 
Wilks, 250 
Willstatter, 240 
Wilson, 375 
Winkler, 156 
Wohler, 273, 375 
Wolffenstein, 431 
Wood, 404 
Wurtz, 124, 125, 4U 



Absolute temperature, 26 
Absorption, 184 
Accumulators, 325 
Acetylene, 297-8 
Acid, antimonic, 401 

arsenic, 395-6 

arsenious, 394 

boracic, 269, 271-2 

boric, 269, 271-2 

bromic, 527 

bromous, 527 

carbamic, 348 

carbonic, 305-6 

Caro's, 478 

chloric, 516-17 

chlorosulphonic, 481 

chlorous, 51516 

chromic, 487, 489 

dichromic, 487 

formic, 299 

graphitic, 287 

hydrazoic, 351 

hydriodic, 531-3 

hydrobromic, 522-6 

hydrochloric, 505-10 

hydrochloroplumbic, 327-8 

hydrochlorostannic, 320 

hydrocyanic, 215 

hydrofluoboric, 271 

hydrofluoric, 494-6 

hydrofluosilicic, 313 

hydrosulphurous, 479 

hypobromous, 526 

hypochlorous, 514r-5 

hypoiodous, 534 

hyponitrous, 369 

Acid, hypophosphorie, 386-7 
hypophosphorous, 385 
iodic, 534 
mellitic, 287 
metaphosphoric, 389-90 
nitric, 352-60 

examples of the action o/, 

structure, 370 
nitrohydroxylaminic, 369 
nitrososulphuric, 470 
nitrous, 367-9, 370 
orthophosphoric, 387-8 
oxalic, 299 
percarbonic, 306 
perchloric, 517-8 
perchromic, 48990 
periodic, 535 
persulphuric, 477-8 
phosphoric, 387-8 
phosphorous, 386 
prussic, 215 
pyrophosphoric, 388-9 
silicic, 311 
stannic, 319 
sulphovinic, 296 
sulphuric, 468-77 

constitution of, 477 

history of, 46&-9 

manufacture of, 469-76 

properties of, 476-7 
sulphurous, 466-7 
tetrathionic, 478 
thionic, 478-9 
thiosulphuric, 478-9 
Affinity, 116-8 



Air, 448-53 

ozone in, 421 
Alabaster, 244 
Alchemy, 4 
Alcohol, 4 
Alkali, 194, 211 
Alkali metals, 191-216 
Alkaline earth metals, 239-52 
Allotropy, 283, 317 
Alum, 273, 280 
Alumina, 273, 277 
Aluminium, 273-7 

compounds, 277-81 
Alunite, 280 
Amalgams, 261-2 
Amethyst, 277 
Ammonia, 337-44, 348-50 
Ammonium, 344-6 

amalgam, 345 

compounds, 346-8 
Amphoteric substances, 278, 420 
Anhydrite, 244 
Anions, 94 
Anode, 93 
Anthracite, 287-9 
Antichlor, 467 
Aiitimoiiiates, 402 
Antimoniurettod hydrogen, 400 
Antimony, 398-400 

compounds, 400-3 
Apatite, 373 
Aqua fortis, 496 
Aqua regia, 496 
Argon, 188-9 
Arsenates, 395, 396 
Arsenic, 390-2 

compounds, 392-7 

tests for, 397 
Arsenites, 394-6 
Arseniuretted hydrogen, 392 
Arsine, 392 
Asbestos, 240 
Atmosphere, 448-53 
Atomic heats, 91 

hydrogen blowpipe, 185 

numbers, 157, 166, 167, 168, 
169, 574 

Theory, 8-17 

weights, 33, 52-9, 574 
Atoms, 10 

Available chlorine, 249 
Avogadro's Hypothesis, 35 

Barium, 251 

compounds, 252 
Basic oxides, 417 
Bauxite, 277 
Bessemer process, 548 
Bismuth, 403-4 

compounds, 4046 
Black ash, 203 
Blacklead, 283 

Bleaching, 466, 467, 504, 505 
Bleaching-powder, 248-50 
Boiling-point method of M.W. 

determination, 1113 
Bone-ash, 373 
Borax, 268, 269, 272 
Borofluorides, 269, 271 
Boron, 268-70 

compounds, 270-2 
Boyle's Law, 26 
Brass, 227 
Briii's process, 252 
Britannia metal, 318 
Bromates, 527 
Bromides, 525 
Bromine, 518-22 

compounds, 522-8 
Bronze, 314 

Cadmia, 258 
Cadmium, 258-9 

compounds, 258 

plating, 259 
Calamine, 254, 258 
Calcite, 244 
Calcium, 243-5 

carbonate, 247 

compounds, 245-50 

hydroxide, 246 

oxide, 245 

phosphate, 248, 373 
Caliche, 209, 528, 529 
Calmia, 258 
Calorie, 41 
Calx, 410-14 
Carbon, 282-94 

allotropy, 283 

compounds, 294-308 

dioxide, 303-7 

disulphide, 308 

monoxide, 299-303 
Carbonado, 284 
Carbonates, 305 



Carbonyls, 301 
Carborundum, 313 
Carnallite, 212, 240 
Cassiterite, 319 
Cast iron, 547 
Catalysis, 140-4 
Cathode, 93 

rays, 159 
Cations, 94 
Caustic potash, 212 
Caustic soda, 196-200 
Cement, 246, 247 
Cementite, 547, 551 
Cerussite, 322 
Chalcocite, 218 
Chalk, 244 
Charcoal, 293 
Charles's Law, 27 
Chlorates, 516-17 
Chlorine, 496-505 

compounds, 50518 
Chlorites, 516 
Chroma tes, 487 
Chrome alum, 486 
Chrome iron ore, 483 
Chromite, 483 
Chromium, 483-4 

compounds, 484-90 

plating, 484 
Cinnabar, 259, 260, 266 
Clarain, 289 
Clark's method of softening water, 


Clay, 274 

Clay ironstone, 554 
Coal, 287-9 
Coarse metal, 219 
Cobalt, 558 

compounds, 55860 
Coke, 290 
Colcmamte, 272 
Colloids, 144-6 
Combustion, 408-14 
Conductivity, 132 
Conservation of Matter, Law of, 


Constant Composition, Law of, 10 
Copper, 217-22 

compounds, 222-7 
Coprolites, 373 
Corrosive sublimate, 264 
Coulomb, 97 

Critical temperature and pressure, 


Crocoisite, 483 
Cryolite, 274 
Crystal structure, 163 
Crystalloids, 144-6 
Cupellation, 228 
Cuprite, 218, 222 
Cyanamide, 248, 342 
Cyanides, 215, 216 
Cycle, Nitrogen, 370-2 

Deacon's process, 142, 499-501 

Deliquescence, 444 

Density of gases, 42, 43 

Dephlogisticated air, 41214 

Dialysis, 144 

Diamond, 283-5 

Diffusion, 74, 75 

Dissociation, 81-7 

Dolomite, 240, 244 

Dulong and Petit's Law, 54, 57, 91, 


Durain, 289 
Duralumin, 277 
Dutch process for white lead, 328, 


Eau de Javelle, 515 

Eau de Labarraque, 515 

Efflorescence, 444 

Effusion, 75 

Electrodes, 93 

Electrolysis, 93-8 

Electrolytes, 94 

Electron, 160 

Electronic structure of atoms, 


Elements, classification of, 149-57 
Emeralds, 277 
Emery, 278 

Endothormic reactions, 87 
Enzymes, 143 
Equations, 60-9 
Equilibrium, 118-25 
Equivalent weights, 1924 
Ethylene, 296, 297 
Evaporation, 75, 76 
Exothermic reactions, 87 

Fehling's solution, 222 
Felspar, 274 



Ferrite, 551 
Firedamp, 294 
Fluorides, 495-6 
Fluorine, 491-4 

compounds, 494-6 
Fluorspar, 492 
Formulae, 60-9 
Freezing-point, depression of, 105- 

Fusain, 289 

Galena, 322 

Gases, absorption of, 184 

densities of, 42, 43 

effect of pressure and tempera- 
ture on, 24-31 

liquefaction of, 7880 
Gay-Lussac's Law, 34 
Glass, 311-12 
Gold, 234-8 

compounds, 238 
Graham's Law, 74 
Gram Molecular Volume, 65 
Gram Molecular Weight, 65 
Graphite, 285-7 

Guldberg and Waage's Law, 120 
Gypsum, 247 

Haber process for ammonia, 341 
Haematite, 545, 555 
Halogens, 491-535 
Hardness of water, 4458 
Heat, atomic, 54, 57, 91, 155 

specific, 54 ; of gases, 42 
Hehum, 187-90 (also Chap. XVI) 
Hess's Law, 89 
Hetero topes, 169 
Hornsilver, 228, 497 
Hydrates, 433-4 
Hydrazine, 337 
Hydrazoic acid, 337, 351 
Hydrocarbons, 283 
Hydrogen, 179-86 

atomic, 184-5 

ion concentration, 138 

peroxide, 431-7 
Hydrolith, 182 
Hydrolysis, 137 
Hydroxylamine, 351 

latrochemistry, 5 

Indicators, theory of, 137, 138 

Inflammable air, 179 
Ink, 554 
Iodine, 528-31 

compounds, 5315 
Ionic theory, 126-39 
lonization, degree of, 1303 
Iron, 543-52 

compounds, 552-7 

pyrites, 545 

Isomorphism, Law of, 58 
Isotopes, 168-75 

Kaolin, 279 
Kelp, 529 
Kibrlt, 454 
Kiesolguhr, 360, 521 
Kinetic Theory, 70-80 
Kipps apparatus, 304 
Krypton, 188, 189 
Kuhl, 391, 398 

Lapis lazuli, 280 

Lead, 321-3 

compounds, 323-30 
desilverization, 228 

Leblanc process, 200-4 

Lopidolite, 192 

Lignite, 288 

Lirno, 245 

Limonite, 545 

Lithium, 191, 192 
compounds, 192 

Litmus, 137, 138 

Magnalium, 241, 277 
Magnesia, 239, 240, 241 
Magnesium, 239-41 

compounds, 2413 
Magnetite, 554 
Malachite, 218 
Maiiganates, 540 
Manganese, 537-9 

compounds, 539-42 
Marble, 244 
Marcasite, 545 
Marsh-gas, 294 
Marsh's test for arsenic, 397 
Mass-Action, Law of, 120 
Masurium, 637 
Matches, 378-9 
Mercury, 259-61 

compounds, 2626 



Metalloids, 149, 150 
Metals, 149-50 
Methane, 294-5 
Methanol, 302 
Methyl alcohol, 302 

magnesium iodide, 294 

orange, 137, 138 
Mica, 274 

Millon's reagent, 264 
Minium, 325 
Mispickel, 391 
Molecular speeds, 74 

weights, 32-51, 104-13 
Molecules, 33 
Mond's process, 561 
Mortar, 246 
Mosaic gold, 321 
Multiple Proportions, Law of, 14 

Nascent hydrogen, 184 
Neon, 188-90 
Nernst lamps, 242 
Nessler's reagent, 266, 350 
Neutralization, 134 
Nickel, 560-2 

compounds, 562 

Niton, 188, 189 (also Chap. XVI) 
Nitrates, 368, 359 
Nitre, 215 

Nitric oxide, 351-2, 363-4 
Nitrites, 215 
Nitrogen, 332-7 

compounds, 337-72 

peroxide, dissociation of, 83, 84, 


Nitrous oxide, 360-3 
Nitroxyls, 367 
Normal solutions, 23 
Normal temperature and pressure, 

Nucleus of atoms, 164-6, 174-7 

Occlusion of hydrogen, 183-4 

Oleum, 474 

Opal, 309 

Orpiment, 390, 396 

Osmotic pressure, 99-114 

Ostwald's Law, 133 

Oxidation, 428-30 

Oxides, classification of, 416-20 

Oximes, 351 

Oxygen, 7, 407-20 

Oxymuriatic acid, 496 
Ozone, 420-8 

in atmosphere, 421 
Ozonides, 424 * 

Palladium, occlusion of hydrogen 
by, 183 

Parkes' process, 229 

Passivity, 484 

Pattinson's process, 229 

Pearl ash, 213 

Peat, 288 

Peligot's salt, 489 

Perchloric anhydride, 513 

Perhydrol, 433 

Periodates, 535 

Periodic System, 58, 149-57 

Permanganates, 540-2 

Permutite, 446-7 

Peroxides, 418, 419 

Pettenkofer's method of deter- 
mining carbon dioxide, 452 

Pharaoh's serpents, 266 

Phase Rule, 146-8 

Phenolphthalein, 137, 138 

Phlogisticated air, 332 

Phlogiston Theory, 6, 7, 410-4 

Phosgene, 307 

Phosphates, 388-90 

Phosphine, 379-81 

Phosphor-bronze, 318 

Phosphorus, 372-8 
compounds, 379-90 

Photography, chemistry of, 233-4 

Pig iron, 546, 547 

Pitchblende, 158 

Plaster of Paris, 247 

Platinized asbestos, 473 

Plumbago, 285 

Portland cement, 247 

Positive rays, 171-4 

Potash, 211 

Potassium, 210-2 
compounds, 2126 
per-rhenate, 537 

Pressure, gas, 26, 71 

law of partial pressures, 31 

Producer gas, 301 

Proton, 175 

Prussian blue, 216 

Purple of Cassius, 237 

Pyrophosphates, 389 



Quantum theory, 91 
Quartz, 309, 310 
Quicklime, 245 

Radioactivity, 158, 159, 1(>3, 1G4 

Radium, 158 

Rate of reaction, 118 

Rays, a-, -, and 7-, 164 

positive, 171-4 

X-, 159-73 
Realgar, 390, 396 
Reciprocal proportions, Law of, 16 
Reduction, 184, 428-30 
Reversible reactions, 11825 
Rhenium, 537 
Rose's metal, 404 
Rouge, jeweller's, 555 
Rydberg' s constant, 166 

Saccharate, 251 
Sal ammoniac, 338 
Sal sedativum, 269 
Saltcake, 202 
Sand, 309 

Scheele's green, 394 
Schweitzer's reagent, 225 
Sheffield plate, 231 
Shells, electronic, 176-7 
Silica, 308-10 
Silicol, 182 
Silicon, 308-10 

compounds, 31014 
Silver, 227-31 

compounds, 2314 
Soda ash, 205 
Sodium, 194-6 

compounds, 196-210 
Solubility product, 134. 133 
Sorption, 184 
Spathic iron ore, 545, 654 
Specific heats of gases, 42 
Spelter, 255 
Spinels, 278 
Stalactites, 448 
Stalagmites, 448 
Steel, 647-51 

varieties of, 550-1 
Stibine, 400 

Strength of Acids and Bases, 138 
Strontium, 250-1 

compounds, 251 
Sulphates, 477 

Sulphides, 464 
Sulphites, 466 
Sulphur, 407, 453-61 
compounds, 46181 
Sulphuretted hydrogen, 4614 
Sulvari, 453 

Talc, 240 

Tartar emetic, 402 
Tautomerism, 436 
Temper of steel, 550 
Thermochemistry, 87-91 
Thionyl chloride, 480 
Thyroxin, 531 
Tin, 314-8 

compounds, 31821 
Tinstone, 315 
Topaz, 277 
Tourmaline, 272 
Transition elements, 543-62 
Transmutation theory, 2 
Turnbull's blue, 216 
Turquoise, 274 
Tutia, 253 

Ultramarine, 280 
Uranium, 158 

Valency, 55, 155 

van der Waal's equation, 77 

van't HofP s coefficient, 127, 131 

Vapour densities, 36, 42-50 

Vapour pressure, 75, 76, 109-13 

Varech, 629 

Verdigris, 222 

Vitrain, 289 

Vitreosil, 310 

Washing soda, 200-7 
Water, 437-48 

hardness of, 445Hf 

natural, 437-8 
Weldon process, 501, 502 
White lead, 328 
Wood's metal, 404 

X-rays, 159, 160-73 
X-ray spectra, 166 
Xenon, 188, 189 

Zarnikh, 391 
Zibaq, 454 
Zinc, 253-6 

compounds, 256-8 



5. Pentoxide 1 g. phosphorus combined with 1-29 g, oxygen. 
Tetroxide 1 g. 1-032 g. 
Trioxide 1 g. -7742 g. 

Weights of oxygen are in ratio 5:4:3. 

6. Stannous 1 g. tin combined with -2695 g. sulphur. 
Stannic 1 g. -5392 g. 

Weights of sulphur in ratio 1 : 2 (approx,). 

7. (i) 1 g. vanadium combined with 1-389 g. chlorine, 
(ii) 1 g- 2-083 g. 

(iii) 1 g. 2-778 g. 

Weights of chlorine in ratio 2:3:4. 

8. Ferrous 1 g. Iron combined with 1-72 g. " sulphate group.*" 
Ferric 1 g. 2-58 g. 

Weights of " sulphate-group " in ratio 2 : 3. 

9. Suboxide 1 g. oxygen combined with 25-9 g. lead. 
Monoxide 1 g. 12-95 g. 
Sesquioxide 1 g. 8-668 g. 
Peroxide 1 g. 6 475 g. 

Weights of lead in ratio 12 : 6 : 4 : 3. 

10. Auroiis 1 g. gold combined with -18 g. chlorine. 

Auric 1 g. -54 g. 

Weights of gold in ratio 1 : 3. 

No. 2. 
No. 3. 



No. 4. 
No. 6. 

64-3 c.e. 


No. 1. 2024-7 c.o.i. 
No. 2. 2-70 litre*. 
No. 8. 651-5 C. 

No. 7. 81-78. 
No. 8. 32-68. 
No. 9. 31-7. 

No. 4. 164 C.C.B. 
No. 5. 21-4 ac. 


No. 10. 24-38. 

No. 11. 5868. 

No. 12. 2006; MI* 


No. 5. CoAsS. 
No. 6. Mg a Si0 4 . 
No. 7. PtN a H,Ci 4 . 

No. 2. 52-3. 

No. 8. CH 4 Br t . 
No. 9. C^TT^NOa 


No. 5. 22,000 calorie*. 


No. 3. 94-6. 


No. 6. 593 mm. 

No. 7. 181} atmospheres. 

No. 8. 4-07 per cent. 

No. 9. 200. 

JSo. 10. 58-2 C., 6-9 atmospheres. 

No. 11. 4 atmospheres. 

No. 12. 
No. 13. 
No. 14. 
No. 15. 




4000 C. 
No. 16. 1 882 gma. 
No. 17. 173-5. 


No. 4. (3-067 gr. mola. acid. 

j 0-067 gr. rnols. alcohol. 

l0'933 gr. mols. ester and watei. 

No. 10. 0-913. 
No. 11. 0-805. 

No. 12. 


0-016 C. 

P. 413. No. 7. 8-93 litres. 

No. 12. 41-5% ozone by volume. 
No. 13. 0-207% ozone by volume. 

P. 426. No. 6. 29 gma. per iitr. 
No. 7. 1-57 
No. 8. 13-87 













O "a> 





^ * 

<q p 


^5 fa 

4^ ^ 





Neodymium . 




Antimony . 




Neon . 




Argon . 












Niobium .